S4: Chemistry

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• UNIT1:STRUCTURE OF AN ATOM AND MASS SPECTRUM

  Key unit competency
  Interpret simple mass spectra and use them to calculate the relative atomic mass 
  (R.A.M) of different elements. 
  Learning objectives 
  By the end of this unit, students should be able to:
  • Outline the discovery of the sub-atomic particles.
  • Compare the properties of sub-atomic particles.
  • Explain what is an isotope of an element.
  • Assess the relationship between the number of protons and the number of 
  electrons.
  • Calculate the mass number knowing the number of protons and the 
  number of neutrons.
  • Understand the meaning of relative atomic mass and relative abundances
  • Calculate the relative atomic mass of an element, given isotopic masses and 
  abundances.
  • Draw and label the mass spectrometer. 
  • Explain the fundamental processes occurring in the functioning of a mass 
  spectrometer.
  • Interpret different mass spectra. 
  • State the uses of the mass spectrometer. 
  

  • Calculate the relative atomic mass of an element, from a mass spectrum. 

  

  2. What do the three diagrams A, B, and C have in common?
  3. Based on your knowledge concerning atomic structure, what do you think 
  that 
  a) the blue spheres represent? b) the red spheres represent?
  Provide explanations.
  4. Using the information obtained in question (3) write the atomic symbol for 
  each of the diagrams.
  5. Are there some other particle(s) missing from the above diagrams? If yes name 
  the particle(s).
  6. What could you obtain if the atom is broken down?
  Each country has its own culture (language, traditions and norms, attitudes and 
  values, etc.). Our culture defines our identity which is unique to each Rwandan 
  citizen and differentiates us from foreigners; if one element of our culture is rejected 
  or disappears, we become a different Rwandan people. When we introduce foreign 
  cultures to replace ours, we can lose our identity. However, some of our cultural 
  elements such as language can be shared with others to build the social relationship. 
  Similarly, in the atom, the number of protons within the nucleus defines the atomic 
  number, which is unique to each chemical element; the atomic number or the 
  number of protons of an atom defines its identity. If a proton is added or removed 
  from an element, it becomes a different element. Electrons around the nucleus can 
  be lost, gained, or shared to create bonds with other atoms in chemical reactions to 
  produce useful substances, but this does not change the identity of the elements 
  involved.
  1.1. Outline of the discovery of the atom's constituents and 
  their properties
  Activity 1.1
  1. Regardless of some exceptions, all atoms are composed of the same components.
   True or False? If this statement is true why do different atoms 
  have different chemical properties?
  2. The contributions of Joseph John Thomson and Ernest Rutherford led the 
  way to today’s understanding of the structure of the atom. What were 
  their contributions? 
  3. Explain the modern view of the structure of the atom?
  4. Using your knowledge about atom, what is the role each particle plays in 
  

  an atom

  1.1.1. Constituents of atoms and their properties
  Atoms are the basic units of elements and compounds. In ordinary chemical 
  reactions, atoms retain their identity. An atom is the smallest identifiable unit of 
  an element. There are about 91 different naturally occurring elements. In addition, 
  scientists have succeeded in making over 20 synthetic elements (elements not 
  found in nature but produced in Laboratories of Reasearch Centers).
  An element is defined as a substance that cannot be broken down by ordinary 
  chemical methods in simpler substances. Some examples of elements include 
  hydrogen (H), helium (He), potassium (K), carbon (C), and mercury (Hg). In an 
  element, all atoms have the same number of protons or electrons but the number 
  of neutrons can vary. A substance made of only one type of atom is also called an 
  element or elemental substance, for example: hydrogen (H2
  ), chlorine (Cl2
  ), sodium 
  (Na). Elements are the basic building blocks of more complex matter. 
  A compound is a matter or substance formed by the combination of two or more 
  different elements in fixed ratios. For example, Hydrogen peroxide (H2
  O2
  ) is a 
  compound composed of two elements, hydrogen and oxygen, in a fixed ratio (2:2).
  During the early twentieth century, scientists discovered that atoms can be divided 
  into more basic particles. Their findings made it clear that atoms contain a central 
  portion called the nucleus. The nucleus contains protons and neutrons. Protons 
  are positively charged, and neutrons are neutral. Whirling around the nucleus are 
  particles called electrons which are negatively charged. The relative masses and 
  

  charges of the three fundamental particles are shown in Table 1.1

  

  The mass of an electron is very small compared with the mass of either a proton or 
  

  a neutron.
  

  The charge on a proton is equal in magnitude, but opposite in sign, to the charge 
  

  on an electron. 
  

  1.1.2. Discovery of the atom constituents.
  

  The oldest description of matter in science was advanced by the Greek philosopher 
  Democritus in 400 BC.
  He suggested that matter can be divided into small particles up to an ultimate 
  particle that cannot be divided, and called that particle atom. Atoms came from the 
  Greek word atomos meaning indivisible.
  The work of Dalton and other scientists such as Avogadro, etc., contributed more 
  so that chemistry was beginning to be understood. They proposed new concepts 
  of atom, and from that moment scientists started to think about the nature of the 
  atom. What are the constituents of an atom, and what are the features that make 
  atoms of the various elements different?
  In 1808 Dalton published his book called A New System of Chemical Philosophy, in 
  which he presented his theory of atoms:
  a) Dalton’s Atomic Theory
  1. Each element is made up of tiny particles called atoms.
  2. The atoms of a given element are identical; the atoms of different elements are 
  different in some fundamental way or ways.
  3. Chemical compounds are formed when atoms of different elements combine with 
  each other. A given compound always has the same relative numbers and types of 
  atoms.
  4. Chemical reactions involve reorganization of the atoms—changes in the way they 
  

  are bound together. The atoms themselves are not changed in a chemical reaction.

  

  Figure 1.1: John Dalton’s Atomic Model

  b)Discovery of Electrons and Thomson’s Atomic Model
  In 1897 J. J. Thomson (1856–1940) and other scientists conducted several 
  experiments, and found that atoms are divisible. They conducted experiments with 
  

  gas discharge tubes. A gas discharge tube is shown in Figure 1.2.

  

  Figure 1.2: Gas discharge tube showing cathode rays originating from the cathode

  The gas discharge tube is an evacuated glass tube and has two electrodes, a cathode 
  (negative electrode) and an anode (positive electrode). The electrodes are connected 
  to a high voltage source. Inside the tube, an electric discharge occurs between the 
  electrodes.
  The discharge or ‘rays’ originate from the cathode and move toward the anode, and 
  hence are called cathode rays. Using luminescent techniques, the cathode rays are 
  made visible and it was found that these rays are deflected away from negatively 
  charged plates. The scientist J. J. Thomson concluded that the cathode ray consists 
  of negatively charged particles, and later they were called electrons.
  Thomson postulated that an atom consisted of a diffuse cloud of positive charge 
  with the negative electrons embedded randomly in it. This model, shown in Figure 
  1.3, is often called the plum pudding model because the electrons are like raisins 
  

  dispersed in a pudding (the positive charge cloud), as in plum pudding.

  

  Figure 1.3.The plum pudding model of the atom

  In 1909 Robert Millikan (1868–1953) conducted the famous charged oil drop 

  experiment and came to several conclusions: He found the magnitude of the charge 

  of an electron to be equal to -1.602 x 10-19c. From the charge-to-mass ratio(e/m) 

  determined by Thomson, the mass of an electron was also calculated.

  

  c)Discovery of Protons and Rutherford’s Atomic Model

  

  

   

  Figure 1.4: A cathode-ray tube with a different design and with a perforated cathode

  The proton was observed by Ernest Rutherford and James Chadwick in 1919 as a 
  

  particle that is emitted by bombardment of certain atoms with α-particles.

  

  Figure 1.5: Rutherford’s experiment on α-particle bombardment of metal foil

  Rutherford reasoned that if Thomson’s model were accurate, the massive α-particles 
  should crash through the thin foil like cannonballs through gauze, as shown in Figure 
  1.6(a). He expected α-particles to travel through the foil with, at the most, very minor 
  deflections in their paths. The results of the experiment were very different from 
  those Rutherford anticipated. Although most of the α- particles passed straight 
  through, many of the particles were deflected at large angles, as shown in Figure 
  1.6(b), and some were reflected, never hitting the detector. This outcome was a great 
  surprise to Rutherford. Rutherford knew from these results that the plum pudding 
  model for the atom could not be correct. The large deflections of the α-particles 
  could be caused only by a center of concentrated positive charge that contains most 
  of the atom’s mass, as illustrated in Figure 1.6(b). Most of the α-particles pass directly 
  through the foil because the atom is mostly an open space. The deflected α-particles 
  are those that had a “close encounter” with the massive positive center of the atom, 
  and the few reflected α-particles are those that made a “direct hit” on the much more 
  massive positive center.
  In Rutherford’s mind these results could be explained only in terms of a nuclear 
  atom—an atom with a dense center of positive charge (the nucleus) with electrons 
  

  moving around the nucleus at a distance that is large relative to the nuclear radius.

  

  d)Discovery of Neutrons
  In spite of the success of Rutherford and his co-workers in explaining atomic 
  structure, one major problem remained unsolved.
  If the hydrogen contains one proton and the helium atom contains two protons, 
  the relative atomic mass of helium should be twice that of hydrogen. However, the 
  relative atomic mass of helium is four and not two. 
  This question was answered by the discovery of James Chadwick, English physicist 
  who showed the origin of the extra mass of helium by bombarding a beryllium foil 
  

  with alpha particles. (See figure 1.7)

  

  Figure 1.7. Chadwick’s experiment

  In the presence of beryllium, the alpha particles are not detected; but they displace 
  uncharged particles from the nuclei of beryllium atoms. These uncharged particles 
  cannot be detected by a charged counter of particles.
  However, those uncharged particles can displace positively charged particles from 
  another substance. They were called neutrons.The mass of the neutron is slightly 
  greater than that of proton.
  Figure 1.8 shows the location of the elementary particles (protons, neutrons, and 
  electrons) in an atom. There are other subatomic particles, but the electron, the 
  proton, and the neutron are the three fundamental components of the atom that 
  are important in chemistry.
  
  

  1.2. Concept of atomic number, mass number, and isotopic 

  mass

  

  1. Compare the two sodium isotopes in the figures above.
  2. From your observation, how do you define the isotopes of an element?
  3. How is the mass number, A, determined? 
  4. What information is provided by the atomic number, Z? 
  5. What is the relationship between the number of protons and the number of electrons 
  in an atom? 
  6. Where are the electrons, protons, and neutrons located in an atom? 
  7. Why is the mass of an atom concentrated in the center?
  8. Sodium-24 and sodium-23 react similarly with other substances. Explain the statement
  9. Say which one(s) of the following statements is(are) correct and which one(s) is(are) 
  wrong: (i) isotopes differ in their number of electrons, (ii) isotopes differ in their mass 
  numbers, (iii) isotopes differ in their number of protons, (iv) isotopes differ by their 
  number of neutrons, (v) all the statements are wrong.
  The atomic number denotes the number of protons in an atom’s nucleus. The mass 
  number denotes the total number of protons and neutrons. Protons and neutrons 
  are often called nucleons. By convention, the atomic number is written on the left 
  side of the element symbol as a subscript, and the mass number on the same side 
  

  but as a superscript. 

  

  Some atoms of the same of element have the same atomic number, but different 
  mass numbers. This means a different number of neutrons. Such atoms are called 
  isotopes of the element.
  Isotopes are atoms of the same element with different masses; they are atoms containing 
  the same number of protons but different numbers of neutrons.
  In a given atom, the number of protons, also called “atomic number” is equal to the 
  number of electrons because the atom is electrically neutral. The sum of the number 
  of protons and neutrons in an atom gives the mass number of that atom.
  Mass number = number of protons + number of neutrons
  

   = atomic number + neutron number

  

  Atomic number, Mass number, protons, Electrons, Isotopes, 
  neutron
  a) The atomic number tells you how many……………………………. and 
  ……………………………………………………. are in an atom.
  ......................................................is the number written as subscript on the left of 
  the atomic symbol.
  b) The total number of protons and neutrons in an atom is called the 
  ……………………………………………………………..
  c) Atoms with the same number of protons but different number of neutrons are called ………………………………………….
  d) The subatomic particle that has no charge is called a 
  

  ………………………………………………

  1.3. Calculation of relative atomic mass of elements with 
  

  isotopes

  

  Relative atomic mass, symbolized as R.A.M or Ar, is defined as the mass of one atom 
  of an element relative to 1/12 of the mass of an atom of carbon-12, which has a 
  mass of 12.00 atomic mass units. The relative atomic mass, also known as the atomic 
  weight or average atomic weight, is the average of the atomic masses of all of the 
  element’s isotopes.
  Relative isotopic mass is like relative atomic mass in that it deals with individual 
  isotopes. The difference is that we are dealing with different forms of the same 
  element but with different masses. 
  Thus, the different isotopic masses of the same elements and the percentage 
  abundance of each isotope of an element must be known in order to accurately 
  calculate the relative atomic mass of an element.
  Notice: Remember that mass number is not the same as the relative atomic mass or 
  isotopic mass! The mass number is the number of protons + neutrons; while relative 
  atomic mass (or isotopic mass) is the mass if you were to somehow weigh it on a 
  

  balance.

  

  By applying the same formula, the relative abundance of the isotopes may be 
  calculated knowing the relative atomic mass of the element and the atomic masses 
  of the respective isotopes.
  Example 2: Chlorine contains two isotopes 35Cl and 37Cl, what is the relative abundance of 
  each isotope in a sample of chlorine if its relative atomic mass is 35.5? 
  

  Solution:

  
  

  
  

  
  

  3. Inlet system is also known as which of the following?
  a) Initial system
  b) Sample reservoir
  c) Sample handling system
  

  d) Element injection system

  The mass spectrometer is an instrument that separates positive gaseous atoms 
  and molecules according to their mass-charge ratio and records the resulting mass 
  spectrum. 
  In the mass spectrometer, atoms and molecules are converted into ions. The ions are 
  separated as a result of the deflection which occurs in the magnetic field.
  The basic components of a mass spectrometer are: vaporisation chamber (to 
  produce gaseous atoms or molecules), ionization chamber (to produce positive 
  ions), accelerating chamber (to accelerate the positive ions to a high and constant 
  velocity), magnetic field (to separate positive ions of different m/z ratios), detector 
  (to detect the number and m/z ratio of the positive ions) and the recorder (to plot 
  

  the mass spectrum of the sample)

  

  A mass spectrometer works in five main stages, namely vaporization, ionization, 
  acceleration, deflection, and detection to produce the mass spectrum.
  Stage 1: Vaporization
  At the beginning the test sample is heated until it becomes vapour and is introduced 
  as a vapour into the ionization chamber. When a sample is a gas, it can directly be 
  introduced into the ionization chamber.
  Stage 2: Ionisation
  The vaporized sample passes into the ionization chamber (with a positive voltage of 
  about 10,000 volts). The electrically heated metal coil gives off electrons which are 
  attracted to the electron trap which is a positively charged plate.
  The particles in the sample (atoms or molecules) are therefore bombarded with a 
  stream of electrons (from the electrons gun), and some of the collisions are energetic 
  enough to knock one or more electrons out of the sample particles to make positive 
  ions. Mass spectrometers always work with positive ions.
  Most of the positive ions formed will carry a charge of +1 because it is much more 
  difficult to remove further electrons from an already positively charged ion.
  Most of the sample molecules are not ionized at all but are continuously drawn off by 
  vacuum pumps which are connected to the ionization chamber (figure 1.9). Some of 
  the molecules are converted to negative ions through the absorption of electrons.
  
  Figure 1.10: A typical ionisation chamber and the nearby accelerating plates
  
  The repeller plate absorbs these negative ions. A small proportion of the positive 
  ions which are formed may have a charge greater than one (a loss of more than one 
  electron). These are accelerated in the same way as the singly charged positive ions.
  Stage 3: Acceleration
  The positive ions are accelerated by an electric field so that they move rapidly 
  through the machine at high and constant velocity. 
  Stage 4: Deflection
  The ions are then deflected by a magnetic field according to their mass to charge 
  ratios. Different ions are deflected by the magnetic field at different extents. The 
  extent to which the beam of ions is deflected depends on four factors:
  1. The magnitude of the accelerating voltage (electric field strength). Higher voltages 
  result in beams of more rapidly moving particles to be deflected less than the 
  beams of the more slowly moving particles produced by lower voltages.
  

  2. Magnetic field strength. Stronger fields deflect a given beam more than weaker 
  fields.
  3. Masses of the particles. Because of their inertia, heavier particles are deflected less 
  than lighter particles that carry the same charge.
  4. Charges on the particles. Particles with higher charges interact more strongly with 
  magnetic fields and are thus deflected more than particles of equal mass with 
  smaller charges
  The two last factors (mass of the ion and charge on the ion) are combined into the 
  mass/charge ratio. Mass/charge ratio is given the symbol m/z (or sometimes m/e).
  For example, if an ion had a mass of 28 and a charge of 1+, its mass/charge ratio 
  would be 28. An ion with a mass of 56 and a charge of 2+ would also have a mass/
  

  charge ratio of 28.

  

  In the figure 1.11 above, ion stream A is most deflected: it will contain ions with the 
  smallest mass/charge ratio. Ion stream C is the least deflected: it contains ions with 
  the greatest mass/charge ratio. Assuming 1+ ions, stream A has the lightest ions, 
  stream B the next lightest and stream C the heaviest. Lighter ions are going to be 
  more deflected than heavy ones.
  Stage 5: Detection
  The beam of ions passing through the machine is detected electrically. As they 
  pass out of the magnetic field, ions are detected by an ion detector which records the 
  position of the ions on the screen and the number of ions that hit the screen at each 
  position. These two pieces of information are used to produce a mass spectrum for 
  the sample.
  A flow of electrons in the wire is detected as an electric current which can be 
  amplified and recorded. The more ions arriving, the greater the current.
  Detecting the other ions
  How might the other ions be detected (those in streams A and C which have been 
  lost in the machine)?
  Remember that stream A was most deflected. To bring them on to the detector, you 
  would need to deflect them less by using a smaller magnetic field.
  To bring those with a larger m/z value (the heavier ions if the charge is +1) to the 
  detector you would have to deflect them more by using a larger magnetic field.
  If you vary the magnetic field, you can bring different ion streams, one at a time 
  on the detector to produce a current which is proportional to the number of ions 
  arriving. The mass of each ion being detected is related to the size of the magnetic 
  field used to bring it on to the detector. The machine can be calibrated to record 
  

   

  Recorder

  The detector of a typical instrument consists of a counter which produces a current 
  that is proportional to the number of ions which strike it. Through the use of electron 
  multiplier circuits, this current can be measured so accurately that the current caused 
  by just one ion striking the detector can be measured. The signal from the detector is 
  fed to a recorder, which produces the mass spectrum. In modern instruments, the 
  output of the detector is fed through an interface to a computer. The computer can 
  store the data, provide the output in both tabular and graphic forms, and compare 
  the data to standard spectra, which are contained in spectra libraries also stored in 
  

  the computer.

  
  

  This is an example a mass spectrum of unknown element that has 2 isotopes.

  Checking up 1.4
  1. Use the list of the words given below to fill in the blank spaces. Each word 
  will be used once.
  Vaporization chamber, mass spectrum, velocity, ionization, deflection, 
  detector, acceleration
  A sample of the element is placed in the _________ where it is converted 
  into gaseous atoms. The gaseous atoms are ionized by bombardment of 
  high energy electrons emitted by a hot cathode to become positive ions (in 
  practice, the voltage in the ________chamber is set in such a way that only one 
  electron is removed from each atom). The positive ions (with different masses) 
  are then given a high and constant _________by two negatively charged 
  plates: the process is called_________. The positive ions are then deflected by 
  the magnet field. This process is called ____________ (ions with smaller mass 
  will be deflected more than the heavier ones). These ions are then detected 
  by the ion _________. The information is fed into a computer which prints out 
  the________ of the element.
  2. The correct order for the basic features of a mass spectrometer is...
  a. acceleration, deflection, detection, ionization
  b. ionisation, acceleration, deflection, detection
  c. acceleration, ionisation, deflection, detection
  d. acceleration, deflection, ionisation, detection
  3. Which one of the following statements about ionisation in a mass spectrometer 
  is incorrect?
  a. gaseous atoms are ionised by bombarding them with high energy electrons
  b. atoms are ionised so they can be accelerated
  c. atoms are ionised so they can be deflected 
  d. it doesn’t matter how much energy you use to ionise the atoms
  4. The path of ions after deflection depends on...
  a. only the mass of the ion
  b. only the charge on the ion
  c. both the charge and the mass of the ion
  d. neither the charge nor the mass of the ion
  

  5. Which of the following species will be deflected to the greatest extent?

  

  

  6. Which of the following separates the ions according to their mass-to-charge 
  ratios?
  a) Ion source
  b) Detector
  c) Magnetic sector
  

  d) Electric sector

  1.5. Interpretation of mass spectra.

  
  

  
  

  The mass spectrum of an element shows how you can find out the masses and 
  relative abundances of the various isotopes of the element and use that information 
  to calculate the relative atomic mass of the element.
  The relative size of the peaks gives you a direct measure of the relative abundances 
  of the isotopes. The tallest peak is often given an arbitrary height of 100 but you 
  may find other scales used; it doesn’t matter. You can find the relative abundances 
  by measuring the lines on the stick diagram.
  

  In this case, the two isotopes (with their relative abundances) are:

  
  

  

  

  

  

  

  a. How many isotopes does magnesium possess
  b. Estimate the isotopic mass of each of the magnesium isotopes
  c. Estimate the relative abundance for each of the isotopes of magnesium
  1.6. Uses of the mass spectrometer and involved calculations
  Activity 1.6
  1. Mass spectrometers are used to determine which of the following?
  a) The atomic mass
  b)Composition in sample
  c) Concentration of elements in sample
  2. The mass spectrum of an element, A, contained four lines at mass/charge 
  of 54; 56; 57 and 58 with relative intensities of 5.84; 91.68; 2.17; 0.31 
  respectively. Explain these data and calculate the relative atomic mass of A
  1.6.1. Calculation of RAM using mass spectrum
  When the mass spectrum of the element is given, you can calculate the relative 
  

  atomic mass of that element by using the information from the mass spectrum.

  

  

  1.6.2. Uses of mass spectrometer
  In addition to the use of mass spectrometer in the determination of isotopes of 
  elements and their relative abundances, other applications of mass spectrometry 
  are:
  • Pharmaceutical: drug discovery, combinatorial chemistry, pharmacokinetics, 
  drug metabolism.
  • Clinical: neonatal screening, haemoglobin analysis, drug testing.
  • Environmental: water quality, soil and groundwater contamination, food 
  contamination, pesticides on foods.
  • Geological: oil composition.
  

  • Biotechnological: the analysis of proteins, peptides

  Checking up 1.6
  1. Which of the following is not done through mass spectrometry?
  a. Calculating the isotopic abundance of elements
  b. Investigating the elemental composition of planets
  c. Confirming the presence of O-H and C=O in organic compounds
  d. Calculating the molecular mass of organic compounds
  2. Mass spectra enable you to find relative abundances of the isotopes of a 
  particular element.
  a) What are isotopes?
  b) Define relative atomic mass.
  c) The mass spectrum of strontium contains the following relative abundances 
  

  for 1+ ions:

  

  a) Explain why there are two separate groups of peaks.
  b) State what causes each of the 5 lines.
  c) Explain the approximate relative heights of the lines at 35 and 37.
  d) Why cannot you predict the relative heights of the two clusters of lines (35/37 
  

  and 70/72/74)?

  1.7. End unit assessment
  I. Multiple choice questions
  1.Which of the following is true regarding a typical atom?
  a. Neutrons and electrons have the same mass.
  b. The mass of neutrons is much less than that of electrons.
  c. Neutrons and protons together make the nucleus electrically neutral.
  d. Protons are more massive than electrons
  2. Which of the following statements is(are) true? For the false statements, 
  correct them.
  a. All particles in the nucleus of an atom are charged.
  b. The atom is best described as a uniform sphere of matter in which 
  electrons are embedded.
  c. The mass of the nucleus is only a very small fraction of the mass of the 
  entire atom.
  d. The volume of the nucleus is only a very small fraction of the total 
  volume of the atom.
  e. The number of neutrons in a neutral atom must equal the number of 
  electrons.
  3. Each of the following statements is true, but Dalton might have had trouble 
  explaining some of them with his atomic theory. Give explanations for the 
  following statements.
  a. Atoms can be broken down into smaller particles.
  b. One sample of lithium hydride is 87.4% lithium by mass, while another 
  sample of lithium hydride is 74.9% lithium by mass. However, the two 
  samples have the same chemical properties
  4. In mass spectrometer, the sample that has to be analysed is bombarded 
  with which of the following?
  a. Protons
  b. Electrons
  c. Neutrons
  

  d. Alpha particles

  5. Mass spectrometer separates ions on the basis of which of the following?
  a. Mass
  b. Charge
  c. Molecular weight
  d. Mass to charge ratio
  6. In a mass spectrometer, the ions are sorted out in which of the following 
  ways?
  a. By accelerating them through electric field
  b. By accelerating them through magnetic field
  c. By accelerating them through electric and magnetic field
  d. By applying a high voltage
  7. The procedure for mass spectroscopy starts with which of the following 
  processes?
  a. The sample is bombarded by electron beam
  b. The sample is accelerated by electric plates
  c. The sample is converted into gaseous state
  d. The ions are detected
  8. Which of the following ions pass through the slit and reach the collecting 
  plate?
  a. Negative ions of all masses
  b. Positive ions of all masses
  c. Negative ions of specific mass
  d. Positive ions of specific mass
  9. Which of the following statements is not true about mass spectrometry?
  a. Impurities of masses different from the one being analysed interferes with 
  the result
  b. It has great sensitivity
  c. It is suitable for data storage
  d. It is suitable for library retrieval
  10. In a mass spectrometer, the sample gas is introduced into the highly 
  evacuated spectrometer tube and it is ionised by the electron beam.
  a. True
  b. False
  II. Short and long answer questions
  11. What are the three fundamental particles from which atoms are built? What 
  are their electric charges? Which of these particles constitute the nucleus of 
  an atom? Which is the least massive particle of the fundamental particles?
  12. Verify that the atomic weight of lithium is 6.94, given the following 
  

  information:

  

  a. Describe the different steps involved in ionizing the particles of a sample
  b. (i) Which two properties of the ions determine how much they are 
  deflected by the magnetic field? What effect does each of these 
  properties have on the extent of deflection?
  (ii) Of the three different ion streams in the previous diagram, why is the 
  ion stream C least deflected?
  (iii) What would you have to do to focus the ion stream C on the detector?
  c. Why is it important that there is a vacuum in the instrument?
  d. Describe briefly how the detector works.
  14. (a) A mass spectrum of a sample of indium shows two peaks at m/z = 113 
  and m/z = 115. The relative atomic mass of indium is 114.5. Calculate the 
  relative abundances of these two isotopes. 
   (b)The mass spectrum of the sample of magnesium contains three peaks with 
  the mass-to-charge ratios and relative intensities shown below
  

  i. Explain why magnesium gives three peaks in mass spectrum?

  

  

  
  

   

• UNIT 2:ELECTRONIC CONFIGURATION OF ATOMS AND IONS

  UNIT 2: ELECTRONIC CONFIGURATION OF ATOMS 
  AND IONS
  Key unit Competence
  To relate Bohr’s atomic model with atomic spectrum of Hydrogen, write electronic 
  configuration of atoms and ions using s, p, d and f atomic orbitals and interpret 
  graphical information related to ionization energy of elements.
  Learning objectives
  By the end of this unit, students should be able to:
  • Explain the stability of atoms using the concept of quantization of energy.
  • Explain the achievements and limitations of Bohr’s atomic model.
  • Explain the existence of energy levels using the data from emission spectra.
  • Describe hydrogen spectral lines and spectral line series
  • Explain the types of spectra in relation with the nature of light
  • Explain the quantum theory of the atom using the quantum numbers.
  • Determine the number and shapes of orbitals in a given energy level or 
  principal quantum number
  • Explain the rules governing the electronic configuration: Aufbau principle 
  and Hund’s rule
  • Explain the relationship between the electronic configuration and the 
  stability of the atoms
  • Interpret the graphs of first ionisation energy against the atomic number.
  

  • Describe the factors which influence the first ionisation energy.

  

  1. What can you see on the image above?
  2. What type of motion is performed by the people on the image?
  3. How does their potential energy change?
  2.1. Bohr’s atomic model and concept of energy levels
  The potential energy of a person walking up ramp increases in a uniform and 
  continuous manner whereas the potential energy of a person walking up steps 
  increases in a stepwise and quantized manner. This can be explained by the values of 
  energy which are continuous for the person walking up ramp while they are discrete 
  (discontinued) for the person walking up steps (Figure 2.1(a) and Figure 2.1(b).
  Niels Bohr (1885-1962) a young Danish physicist working in Rutherford’s laboratory, 
  suggested a model for the hydrogen atom and predicted the existence of line 
  spectra. In his model, based on Planck’s and Einstein’s ideas about quantized energy, 
  Bohr proposed three postulates:
  • An electron can rotate around the nucleus in certain fixed orbits of definite 
  energy without emission of any radiant energy. Such orbits are called 
  

  stationary orbits. 

  

  • An atom can make a transition from its stationary state of higher energy 
  E2
   to a state of lower energy E1 and emit a single photon of frequency ν. 
  Conversely, an atom can absorb an energy at the lower level E1
   and transit 
  to the higher energy level E2
  . That is, the change in energy for a system, 
  ,which can be represented by the equation: 
   where n is an integer (1,2,3,…) and h is Planck’s constant 
  determined from experiment and has a value of 6.626x10-34J.s; ⱱ is the frequency of 
  the electromagnetic radiation absorbed or emitted. Each of these small “packets” 
  of energy is called a photon or a quantum of energy. Energy can be gained or lost 
  only in whole-number multiples of the quantity hν ,
  That is, the change in energy for a system can be represented by the equation:
  , where n is an integer (1,2,3,…).
  An electron does not release energy when it is in its stationary orbit. That is, the 
  electron does not change energy while it is moving around on a given orbit. 
  When an electron on a given orbit absorbs an appropriate quantum of energy, it 
  jumps, it is promoted to a higher energy level; this process is called “excitation” of 
  electron. On the contrary, if the electron loses an appropriate quantum of energy, 
  it falls on the lower energy level by emission of a light corresponding to the lost 
  quantum of energy and the process is called “de-excitation” of electron. As there 
  are many energy levels electrons can be excited to and de-excited from, an atom 
  will have many lines of absorption, each corresponding to a quantum of energy 
  absorbed: this appears as a series of lines called absorption spectrum. In the same 
  way the series of emission lines will produce an emission spectrum( see figure 2.3.a)
  2.1.1. Achievements of Bohr’s Atomic Model
  • Explanation of the stability of an atom
  Based on Rutherford’s atomic model, the electrons move around the nucleus in 
  circular paths called orbits. According to the classical theory of electromagnetism, 
  a charged particle revolving around a charged nucleus would release energy and 
  end up by spiraling into the nucleus; thus the atoms would be unstable. The Bohr’s 
  atomic model makes an assumption of discrete orbit (allowed orbits) to explain 
  why an atom is stable; by doing so, Bohr introduces the concept of quantization of 
  energy. 
  • Explanation of the production of the absorption and emission spectra
  The Bohr’s atomic model explains the origin of atomic absorption and emission 
  spectra.
  2.1.2. Limitations of Bohr Model
  1.Bohr’s theory fails to explain the origin of the spectral lines of multi-electron 
  atoms. 
  It only explains the origin of the spectrum of hydrogen-like species having only one 
  electron such as H, He+, Li2+, Be2+, ........
  The model fails to explain the spectral lines of atoms or species with more than one 
  electron.
  2. According to Bohr, the circular orbits in which electrons revolve are planar. 
  However, modern research has shown that an electron moves around the nucleus in 
  the three dimensional space.
  3. Bohr’s theory fails to account for Zeeman Effect and Stark Effect. Zeeman Effect 
  is the splitting of the spectral lines into thinner and closely- spaced lines when an 
  excited atom is placed in a magnetic field. Stark Effect consists of the splitting of the 
  spectral lines into thinner and closely-spaced lines in presence of electric field.
  4. Bohr’s theory is in contradiction with Heisenberg’s uncertainty principle. Bohr 
  assumes that the electron revolves around the nucleus in circular orbits at fixed 
  distance from the nucleus and with a fixed velocity. However, according to W. 
  Heisenberg, it is not possible to know simultaneously the accurate position and the 
  velocity of a very small moving particle such as an electron. 
  Checking up 2.1
  1. Find out two more examples that you can use to illustrating the concept of 
  quantization.
  2. Discuss the main weakness of Rutherford’s nuclear atom. 
  

  2.2. Hydrogen spectrum and spectral lines

  Activity 2.2

  
  

  Bohr’s atomic model allows to explain the emission spectra of atoms. This happens 
  when excited electrons lose energy in form of electromagnetic radiation and fall to 
  

  lower energy levels. 

  The wave-particle nature of the light

  Light as a wave

  The light is a wave-like phenomenon as shown in Figure 2.2.

  It is characterized by its wavelengths, generally symbolized by the Greek letter 

  lambda, λ, and its frequency, represented by the Greek letter nu1

  , ν. 

  As shown in the Figure below, the wavelength represents the distance between two 

  successive summits/peaks (or two successive troughs).The frequency represents the 

  number of complete wavelengths made by the light per second, also called cycles 

  per second.

  Visible light is composed by different visible lights with different λ and ν.

  But all those lights have the same speed, the speed of light, which, in a vacuum, is 

  1 The letter gamma, γ, may also be used

  
  

  
  

  (Source: http://psychelic-information-theory.com/em_spectrum)
  When an electron is excited or de-excited, the energy absorbed or emitted 
  corresponds to the difference of energy, ΔE, between the final energy level of the 
  electron, E2, and the starting energy level of the electron, E1: E2 – E1 = ΔE = hν. ΔE is 
  positive when E2>E1, this is the case of absorption and excitation of electron; on the 
  other hand ΔE may be negative when E2<E1, in case of emission and de-excitation 
  of electron.
  Figure 2.4 below shows the different series of emission spectra of hydrogen. As 
  you can see, the difference between those series is the final energy level where the 
  electron fall after de-excitation.
  The series have been named according to the scientists who discovered them.
  Ionization of an atom or loss of an electron corresponds to excitation of an electron 
  

  to the level n=∞.

  Figure 2.3.a: continious, absorption and emission spectra
  
  
  

  
  

  
  

  
  

  
  

  Checking2.2
  1. What is the meaning of infinity level in the hydrogen spectral lines?
  2. Given a transition of an electron from n=5 to n=2. Calculate
  i) energy 
  ii) Frequency
  

  iii) Wavelength

  2.3. Atomic spectra

  Activity 2.3

  
  

  Observe the picture above, discuss in groups and answer the following 

  questions.

  a. What do you see on the above photo? 
  b. State the physical phenomenon which is related to the above photo.
  c. Think of any other means of producing the same pattern. List two of them.
  d. What property can you attribute to light with reference to the above process?
  The atomic spectrum is the range of characteristic frequencies or electromagnetic 
  radiations that are readily absorbed or emitted by an atom. It is also known as a line 
  spectrum.
  When white light is passed through a prism, we see a myriad of colors – specifically 
  what we call a rainbow. This dispersion of white light demonstrates that white light 
  contains all the wavelengths of radiations and is thus considered to be continuous. 
  Each color blends into the next with no discontinuity and we get continuous or light 
  spectrum. 
  When elements are vaporized and then thermally excited, they emit light. However, 
  this light is not in the form of a continuous spectrum as observed with white light. 
  Instead, a discrete line spectrum is seen. A series of fine lines of different colors 
  separated by large black spaces is observed. The wavelengths of those lines are 
  characteristic of the element producing them – thus, elements can be identified 
  based on the spectral line data that they produce. 
   Two types of atomic spectra are known: Emission and absorption spectra.
  1. Emission is the ability of a substance to give off light when it interacts with 
  heat. 
  2. Absorption is the opposite of emission, where energy, light or radiation is 
  absorbed by the electrons of a matter.
  NB: A combination of the emission and absorption spectra of a given atom gives a 
  

  continuous spectra.

  

  Checking up 2.3
  1. Different metals, when exposed to a flame, emit different flame colors. 
  Explain the origin of that difference.
  2. Would you expect to see the emission of lines and the absorption lines 
  of a given element to appear at the same place on a photographic plate 
  or not. Explain your answer.
  

  3. How do you explain the many spectral lines for the same element?

  2.4 Orbitals and Quantum Numbers

  Activity 2.4

  1. a)Write the electronic configuration of aluminium atom(Z=13)
  b) Indicate the number of electrons in each energy level/quantum shell
  c) The shells are numbered from inside-outward starting from 1, 2, 3, 4 … 
  which other name is given to these shells?
  d) How did you obtain the exact number of electrons in each energy level/
  

  quantum shell in (c) above?

  To answer the questions that could not be answered by Bohr’s atomic model, other 
  atomic models were proposed. One of those models is the Quantum model that 
  has been developed by the Austrarian physicist Erwin Schrödinger (1887-1961). 
  The model is based on a mathematical equation called Schrödinger equation. This 
  model is based on the following assumptions or hypotheses:
  • An electron moves around the nucleus continuously. However, it is not 
  possible to determine its precise position and velocity at the same time. We 
  can only determine the region, around the nucleus, where there is a high 
  probability of finding that electron at a given time. 
  • The region where the probability of finding electron is high, at more than 
  95%, is called “orbital”; in other words, the orbital is the volume or the 
  space (tridimensional) around the nucleus where there is a high probability 
  of finding the electron.
  Without going into the mathematical development of the Schrödinger equation, we 
  can say that the energy of the electron depends on the orbital where it is located. 
  And an atomic orbital is described by a certain number of “quantum numbers”
  according to the solution of Schrodinger equation, i.e. 3 whole numbers:
  1) The principal quantum number is a positive integer which varies from 1 
  to ∞. The principal quantum number indicates the energy level in an atom where 
  electrons can be located: the higher the n value, the higher the energy level. An 
  electron in energy level n=1 has lowest energy in an atom. The principal quantum 
  number, n, has been traditionally given names with the letters: K(n=1), L(n=2), 
  

  M(n=3), N(n=4), O(n=5), P(n=6)

  

  2)The angular momentum quantum number (l)
  The second quantum number is the angular momentum quantum number 
  represented by the letter, l: it is an integer which can take any value from zero or 
  higher but less than n-1, i.e. equal to: 0,1, 2, 3,….up to n-1. For example if n= 1, l 
  is equal to 0, if n= 2, l can be 0, 1. It is also called secondary or azimuthal quantum 
  number. It indicates the shape of the orbital and is sometimes called the orbital 
  shape quantum number. By tradition, those different shapes of orbitals have been 
  given names or letter symbols: l = 0 = s, l =1 = p, l = 2 = d, l=3 = f 
  3) Magnetic quantum number (ml)
  The magnetic quantum number describes the orientation of the orbital. It is an 
  integer that varies from -l to +l. For example if: l = 0, ml
   can only be 0; if l = 1, ml = 
  -1, 0, +1; if l=2, ml
   = -2, -1, 0, 1, 2. As you can see for each value of l there are (2l+ 1) 
  values of ml corresponding to (2l + 1) orientations under the influence of magnetic 
  field. The s orbital where l is zero and ml has no orientation; it has the shape of a 
  

  sphere as shown in figure 2.4 

  

  

  

  4) The spin quantum number (S)
  The fourth quantum number is the spin quantum number, represented by the 
  symbol S (or ms
   in some books). The electron behaves as a spinning magnet.The 
  spin quantum number is the property of the electron, not the orbital.
  This number describes the spinning direction of the electron in a magnetic field. 
  The direction could be either clockwise or counterclockwise. The electron behaves 
  as if it were spinning about its axis, thereby generating a magnetic field whose 
  direction depends on the direction of the spin. The two directions for the magnetic 
  field correspond to the two possible values for the spin quantum number, S (ms). 
  

  Only two values are possible: s = +1/2 and -1/2 as shown in the Figure 2.7 below.

  

  

  

  

  

  

  

  

  

  Orbital box representation
  An orbital box representation consists of a box for each orbital in a given energy 
  level, grouped by sublevel, with an arrow indicating an electron and its spin.
  Note that two electrons in the same orbital have necessarily opposite spins as 
  indicated in the examples below.
  The table 2.4 shows the electronic configuration of some elements using orbital box 
  representation and applying Hund’s rule.
  N.B: An orbital box representation doesn’t show the real form of the orbital; the forms of the 
  

  different orbitals are shown in Figures 2.4, 2.5 and 2.6.

  

  

  Noble Gas Notation
  All noble gases have completely filled subshells and can be used as a shorthand way 
  of writing electron configurations for subsequent atoms.
  When using this method, the following steps are respected.
  a. Identify the noble gas whose electronic configuration is included in that of 
  the concerned element.
  b.Write the chemical symbol of the identified noble gas within square brackets. 
  We call this the noble gas core.
  c. Add electrons beyond the noble gas core. Note that electrons that are added to the electronic level of the highest principal quantum number (the 
  

  outermost level or valence shell) are called valence electrons.

  
  

  2.6. Relationship between ionization energy, energy levels 

  and factors influencing ionization energy

  
  

  2.6.1. Concept of Ionization energy

  The ionization energy is a measure of the energy needed for an atom, in gaseous 
  state, to lose an electron and become positive ion.
  The first ionisation energy is the energy required to remove one mole of electrons 
  from one mole of atoms in their gaseous state. The example below shows how to 
  

  represent the successive ionization energies of an atom M.

  
  

  
  

  b) To remove the second electron needs greater energy because this electron is closer 
  to the nucleus in a 2p orbital. There is a steady increase in energy required as electrons are removed from 2p and then 2s-orbitals.
  c) The removal of the tenth and eleventh electrons requires much greater amounts 
  of energy, because these electrons are closer to the nucleus in 1s orbital.
  2.6.3. Factors influencing the extent of ionization energy
  The ionization energy is a physical property of elements that can be influenced by 
  some factors:
  1) Size of atom
  The atomic size is the distance between the nucleus and valence shell. As the 
  number of energy levels (shells) increases, the force of attraction between nucleus 
  and valence electron decreases. Therefore, the valence electrons are loosely held 
  by the nucleus and lower energy is required to remove them, i.e. ionization energy 
  decreases with an increase in atomic size and vice versa.This is what happens when 
  you go down a Group.
  2) Nuclear charge
  The nuclear charge is the total charge of all the protons in the nucleus. As the nuclear 
  charge increases, the force of attraction between nucleus and valence electrons 
  increases and hence makes it difficult to remove an electron from the valence 
  shell. The higher the nuclear charge, the higher the ionization energy. This is what 
  happens when you cross a period from left to right.
  3) Screening effect or Shielding effect
  The Screening effect or Shielding effect is due to the presence of inner electrons 
  which have a screening or shielding effect against the attraction of the nucleus 
  towards the outermost electrons. The electrons present in inner shells between 
  the nucleus and the valence shell reduce the attraction between nucleus and the 
  outermost electrons. This shielding effect increases with the increasing number 
  of inner electrons. A strong Shielding effect makes it easier to remove an external 
  electron and hence lowers the ionisation energy.
  2.6.4. Importance of ionization energy in the determination of the 
  chemistry of an element
  Ionization energy provides a basis to understand the chemistry of an element. The 
  following information is provided.
  Determination of metallic or non- metallic character.
  The I.E informs us how the atom will behave chemically: a low I.E indicates that the
  element behaves as metal whereas a high I.E indicates that the element behaves as 
  non-metal.
  The first ionization energies of metals are all nearly below 800kJ mol-1 while those of 
  non- metals are all generally above 800 kJ mol-1.
  Down the group ionization energies decrease so that the elements became more 
  metallic. In groups 14 and 15 there is change from non metallic to metallic character. 
  Across a period from left to right 1st I.E. increases. The elements become less metallic 
  to non- metallic
  
  2.7. End unit assessment
  1. Which of the following is the correct representation of the ground-state 
  electron configuration of molybdenum? Explain what is wrong with each 
  

  of the others

  

  

  

  12. Four possible electron configurations for a nitrogen atom are shown below, but only one represents the correct configuration for a nitrogen atom 
  in its ground state. Which one is the correct electron configuration? Which 
  configurations violate the Pauli Exclusion Principle? Which configurations 
  

  violate Hund’s rule?

  

  a) What factors determine the magnitude of the first ionization energy?
  b. To which group does element W belong? Explain
  c. Would you expect W to be a metal or a non-metal? Explain your answer.
  

  d. Write an equation representing the second ionization of W. 

  
  

  
  

   

  
  

• UNIT 3:FORMATION OF IONIC AND METALLIC BONDS

  UNIT 3: FORMATION OF IONIC AND METALLIC 
  BONDS
  Key unit competence
  Describe how properties of ionic compounds and metals are related to the nature 
  of their bonding
  Learning objectives
  By the end of this unit, student should be able to:
  • Explain why atoms bond together;
  • Explain the mechanisms by which atoms of different elements attain 
  stability; 
  • Explain the formation of ionic bonds using different examples; 
  • Represent ionic bonding by dot-and-cross diagrams; 
  • Describe the properties of ionic compounds based on observations; 
  • Perform experiments to show properties of ionic compounds; 
  • Assemble experimental set up appropriately and carefully;
  • State the factors that influence the magnitude of lattice energy ;
  • Relate the lattice structure of metals to their physical properties;
  • Describe the formation of metallic bonds; 
  • State the physical properties of metals and forces of attraction that hold 
  

  atoms of metal

  Introductory Activity
  1. Look at the pictures below and answer the following questions. Record 
  your answers. 
  a. Observe carefully pictures A, B and C and suggest the similarity between 
  them.
  b. What can you say about the arrangement of chloride and sodium ions in 
  

  the pictures below? c. What holds the chloride and sodium ions together?

  

  People like to bond with each other for many reasons such as: to unite their forces 
  and be stronger, to exchange ideas and produce big things, to found a family, etc. We 
  cannot live in isolation. This inseparability of people can result into strong or weak 
  connection. Similarly, atoms can bond together to form strong or weak connections. 
  Some atoms may not need to bond with others; they are self-sufficient as some 
  people, a small number, may be self-sufficient.
  Connections between atoms are called chemical bonds. Solids are one of the three 
  fundamental states of matter. In molecules, atoms or ions are held together by 
  forces called chemical bonds.There are 3 types of chemical bonds: Ionic, Covalent 
  and Metallic bonds. 
  The type of a bond in molecules is determined by the nature and properties of 
  the bonding atoms. However, in this unit we will only focus on ionic and metallic 
  bonding.
  

  3.1. Stability of atoms and why they bind together

  Activity 3.1
  1. In pairs discuss and write electronic configuration of sodium , neon, argon, 
  magnesium, aluminium, oxygen and chlorine
  2. What happens when oxygen and chlorine gain electrons?
  3. What happens when sodium, magnesium and aluminium lose electrons?
  4. Discuss how atoms of elements can gain their stabilities by either losing or 
  gaining electron(s) on the valence shells and show with evidence that an atom 
  is stable.
  5. How does the formation of an ionic bond between sodium and chlorine reflect 
  the octet rule?
  Like people always relate and connect to others depending on their values, 
  interests, and goals, so do the unstable atoms combine to achieve stability. 
  We know that noble gases are the most stable and unreactive elements in the 
  periodic table. They do not tend to form compounds or combine to themselves
  What do the noble gases have in common? They have a filled outer electron energy 
  level. When an atom loses, gains, or shares electrons through bonding to achieve 
  a filled outermost energy level, the resulting compound is often more stable than 
  individual separate atoms. Neutral sodium has one valence electron. When it gives 
  this electron to chlorine, the resulting Na+ cation has an outermost energy level that 
  

  contains eight electrons

  

  It is isoelectronic (same electronic configuration) with the noble gas neon. On the 
  other hand, chlorine has an outer electron energy level that contains seven electrons. 
  When chlorine gains sodium’s electron, it becomes an anion that is isoelectronic 
  with the noble gas argon.
  The following are examples of how magnesium bonds with oxygen and calcium 
  

  with chlorine:

  
  

  
  

  3.2. Ionic bonding
  Atoms have many ways of combining together to achieve the octet structure, and 
  one of them is the formation of an ionic bond.
  In an ionic bond, electrons are transferred from one atom to another so that they form 
  oppositely charged ions; in other words, one atom loses electron(s) andthe other 
  gains electrons(s).The resulting strong force of attraction between the oppositely
  

  charged ions is what holds them together. Ionic bonding is the electrostatic 

  attraction between positive and negative ions in an ionic crystal lattice.

  3.2.1. Formation of ionic bond
  Activity 3.2
  Draw diagrams to illustrate the formation of ionic compounds in magnesium 
  oxide, magnesium chloride, sodium peroxide, and sodium sulphide.
  The transfer of electrons from one atom to another followed by attraction 
  between positive and negative ions is called ionic bonding. This type of bonding 
  occurs between metals and non-metals. The compounds formed are called ionic 
  compounds. As stated previously, metals try to lose their outer electrons while non 
  metals look to gain electrons to obtain a full outer shell. When metals lose their 
  outer electrons they form positively charged ions called cations. When non-metals 
  gain electrons they form negatively charged ions called anions. An example is 
  

  shown below:

  

  

  

  Checking Up 3.2
  1. For each of the following ionic bonds: Sodium + Chlorine, Magnesium + Iodine, 
  Sodium + Oxygen, Calcium + Chlorine and Aluminium + Chlorine
  a. Write the symbols for each element.
  b. Draw a Lewis dot structure for the valence shell of each element.
  c. Draw an arrow (or more if needed) to show the transfer of electrons to the 
  new element.
  d. Write the resulting chemical formula.
  e. Write the electron configurations for each ion that is formed. Ex. H1+ = 1s0
  2. Solid sodium chloride and solid magnesium oxide are both held together by 
  ionic (electrovalent) bonds. 
  

  a. Using s,p and d notation write down the symbol for and the electronic 
  

  configuration of (i) a sodium ion; (ii) a chloride ion; (iii) a magnesium ion; (iv) 
  an oxide ion. 
  b. Explain what holds sodium and chloride ions together in the solid 
  crystal
  c. Sodium chloride melts at 1074 K; magnesium oxide melts at 3125 K. 
  Both have identical structures. Why is there such a difference in their 
  

  melting points?

  3.2.2. Physical properties of ionic compounds

  Activity 3.3(a)

  Determination of Relative Melting Point of different substances
  Procedure:
  1. Cut a square of aluminum foil that is about5 by 5 cm
  2. Set up a ring stand with an iron ring attached.
  3. Place the aluminum square on the iron ring, as shown at right in Figure 3.6
  4. Obtain a small pea-sized sample of NaCl. Place the sample on the aluminum 
  foil, about 5cm from the center of the square.
  5. Obtain a small pea-sized sample of table sugar. Place the sample on the 
  aluminum foil, about 1 cm from the center of the square, but in the opposite 
  direction from the salt.
  6. Your square of aluminum foil should look like in Figure 3.7.
  7. Light the Bunsen burner and adjust the flame height so that the tip of the 
  flame is just a cm or so below the height of the aluminum foil. 
  8. Observe as the two compounds heat up. 
  9. Set up another sheet of aluminum foil and determine the relative melting 
  points (low vs. high) of the four unknowns. 
  10. Record your results in the table 3.1 below
  

  Caution: if the compounds burn with sparks do not panic.

  

  Conclusions:
  The melting points of ionic compounds are higher than those of covalent 
  compounds; this is due to strong electrostatic forces between opposite charges in 
  the ionic substances compared to the week forces of attraction between molecules 
  in covalent substances . This also explains why all ionic compounds are solid at room 
  

  temperature

  Activity 3.3(b)
  Conductivity in Solution 
  Procedure:
  1. Dissolve a spoonful of NaCl in water. 
  2. Connect the apparatus as shown in figure 3.8 
  3. Make an observation and record your results as in table 3.2 below
  4. Repeat the procedure 1 to 3 above using sugar solution, ethanol and 
  copper(II) sulfate solution
  

  5. Record your results in the table below.

  

  Study questions: 
  1. Give reasons for your observations above.
  2. Solid sodium chloride does not conduct electricity whereas an aqueous 
  solution of sodium chloride does. Explain
  Conclusion: 
  Based on our tests with salt and sugar, the ability to conduct electricity in solution of 
  ionic compounds is much higher than in covalent compounds.
  Activity 3.3(c) Solubility test
  Procedure: 
  1. Using forceps, place 5-8 crystals of each of sodium chloride, magnesium 
  chloride, copper sulphate, calcium carbonate, copper carbonate, sodium 
  sulphate (a small pinch) of the compound into one of the test tubes in test 
  tube rack. 
  2. Half-fill the test tube with distilled water and stir with a clean stirring rod.
  3. Observe if the crystals dissolve in water.
  

  4. Record your findings in a suitable table.

  

  

  Conclusion:
  Water is a good solvent for many ionic compounds but not a solvent for covalent 
  compounds, apart from few exceptions (you will learn about later on).
  Shattering: Why are Ionic compounds generally hard, but brittle?  
  It takes a large amount of mechanical force, such as striking a crystal with a hammer, 
  to force one layer of ions to shift relative to its neighbour.  However, when that 
  happens, it brings ions of the same charge next to each other (Figure3.9). The 
  repulsive forces between like-charged ions cause the crystal to shatter.  When an 
  ionic crystal breaks, it tends to do so along smooth planes because of the regular 
  

  arrangement of the ions.

  

  Checking Up 3.3
  1. The diagrams below show the electric conductivity of distilled water, solid 
  sodium chloride and a solution of sodium chloride respectively. Use the 
  diagrams to explain the observations from the set up.
  i) no light is given out by bulb in A
  ii) no light is given out by bulb in B
  iii) light is given out in C
  
  2.Why are ionic compounds brittle?
  3.Why do ionic compounds have high melting points?
  4.What happens when an electric current is passed through a solution of an 
  

  ionic compound?

  3.2.3. Lattice energy

  Activity 3.4 
  By using information in this student’s chemistry book and other books from the 
  school library, attempt to answer the following questions.
  1. Define lattice energy
  

  2. Explain how the lattice energy is used to describe high melting points of 
  ionic compounds.
  3. What is the bonding force present in ionic compounds?
  4. Why is the melting temperature of magnesium oxide higher than that of 
  magnesium chloride, even though both are almost 100% ionic? 
  5. How is lattice energy of ionic compounds related to their high melting 
  points?
  It is a type of potential energy that may be defined in two ways. In one definition, 
  the lattice energy is the energy required to break apart an ionic solid and convert 
  its component ions into gaseous ions (Endothermic process). On the other hand 
  lattice energy is the energy released when gaseous ions bind to form an ionic solid 
  (Exothermic process). Its values are usually expressed with the units’ kJ/mol.
  Lattice Energy is used to explain the stability of ionic solids. Some might expect 
  such an ordered structure to be less stable because the entropy of the system would 
  be low. However, the crystalline structure allows each ion to interact with multiple 
  oppositely charge ions, which causes a highly favourable change in the enthalpy of 
  the system. A lot of energy is released as the oppositely charged ions interact. It is this 
  that causes ionic solids to have such high melting and boiling points. Some require 
  such high temperatures that they decompose before they can reach a melting and/
  

  or boiling point.

  

  There are two main factors that affect lattice enthalpy.
  a) The charges on the ions
  Sodium chloride and magnesium oxide have exactly the same arrangements of ions 
  

  in the crystal lattice, but the lattice enthalpies are very different.

  

  From the above diagram the lattice enthalpy of magnesium oxide is much greater 
  than that of sodium chloride. This is because in magnesium oxide, +2 ions are 
  attracting -2 ions; in sodium chloride, the attraction is only between +1 and - 1ions.
  b. The radius or the size (volume) of the ions
  The lattice enthalpy of magnesium oxide is also increased relative to sodium chloride 
  because magnesium ions are smaller than sodium ions, and oxide ions are smaller 
  than chloride ions.
  It means that the ions are closer together in the lattice, and that increases the 
  strength of the attractions.
  For example, as you go down Group 17 of the Periodic Table from fluorine to iodine, 
  you would expect the lattice enthalpies of their sodium salts to fall as the negative 
  ions get bigger - and that is the case:
  
  

  3.3. Formation of metallic bonds and physical properties of 

  metals

  
  

  The figure above shows materials commonly used at home. If you reflect back 
  around your house/home you will see hundreds of objects made from different 
  kinds of materials. 
  1. Observe the objects (in picture) and classify them according to the 
  materials they are made of.
  2. Have you ever wondered why the manufacturers choose the material 
  they did for each item?
  3. Why are frying saucepans made of metals and dishes, cups and plates 
  often made of glass and ceramic?
  4. Could dishes be made of metal? And saucepans made of ceramic and 
  

  glass

  3.3.1. Formation of metallic bond
  Another way of combining is the combination between metal atoms to form metallic 
  bond.
  When metal atoms combine together, there is no transfer of electrons since the 
  combining atoms are of the same nature, i.e. all are metals and no one is ready to 
  give up or to capture electrons.
  
  In metallic bonding, all metal atoms put together their valence electrons in a kind of 
  pool of electrons where positive metallic cations seem to bathe. This model is called 
  “Elecron Sea Model” (Fig. 3.13)
  
  Metals have a sea of delocalized electrons within their structure. These electrons 
  have become detached and the remaining atoms have a positive charge. This 
  positive charged is attracted to the delocalized sea of electrons due to electrostatic 
  forces of attraction (forces which result from unlike charges), and as a result has a 
  strong interaction. It is this interaction which makes the metals so hard and rigid. 
  

  Figures 3.13 and 3.14 are representations of metallic bond.

  

  3.3.2. Physical properties of metals
  Activity 3.7: Looking at metals
  1. Collect a number of metal items from your home or school.
  Some examples are listed below: hammer, electrical wiring, cooking pots, jewellery, 
  burglar bars and coins, nails, 
  2. What is the function of each of these objects?
  3. Discuss why you think metal was used to make each object. You should 
  

  consider the properties of metals when you are answering this question

  a. Electrical conductivity
  Activity 3.8 
  Procedure 
  1. Take a dry cell/battery, a torch bulb/ bulb, connecting wires, crocodile clips 
  and connect them. As in the figure 3.17 
  2. Repeat the experiment above using different metals
  

  3. Record your results in a suitable table.

  

  Study questions: 
  1. Compare the relative conductivity of the metals used in the above experiment.
  2. Suggest the purpose of the resistor in the experimental set up.
  Due to the mobile valence electrons of metals, electricity can pass through the 
  metals easily. So they are conductors of electricity. Silver and copper are the best 
  

  conductors of electricity

  Note: mercury is a poor conductor of electricity.

  Thermal conductivity
   
  

  

  Procedure: 
  1. Pour boiling water into the two cups so that they are about half full. 
  2. At the same time, place a metal spoon into one cup and a plastic spoon in 
  the other.
  3. Note which spoon heats up more quickly.
  4. Record your observations.
  Study questions:
  1. Which one heats faster plastic spoon or metallic spoon and why?
  2. Why do we use plastic cups?
  3. Why are cooking pots made of metallic materials not plastics?
  Results: The metal spoon heats up more quickly than the plastic spoon. In other words, the 
  metal conducts heat well, but the plastic does not. 
  Conclusion: Metals are good thermal conductors, while plastic is a poor thermal conductor. 
  The reason is due to the mobility of electrons with transfer of kinetic energy between 
  electrons. This explains why cooking pots are metallic, but their handles are often plastic or 
  wooden. The pot itself must be metal so that heat from the cooking surface can heat up the 
  pot to cook the food inside it, but the handle is made from a poor thermal conductor so that 
  

  the heat does not burn the hand of the person who is cooking.

  c. Malleability and ductility
  Activity 3.10
  Experiments to demonstrate the malleability and ductility of metals
  Materials: wires, nails, hammer, piece of cloth.
  Procedure:
  1. Wrap the material to be test in a heavy plastic or cloth to avoid pieces flying 
  from the material.
  2. Place the material on a flat hard surface
  3. Use a harmer to pound the material flat
  4. Record your observations as malleable or non-malleable.
  Metals can have their shapes changed relatively easily in two different ways i.e.
  Malleable: can be hammered into sheets or 
  Ductile: can be drawn into rods and wires 
  As the metal is beaten into another shape the delocalised electron cloud continues 
  

  to bind the “ions” together.

  

  d) Metal appears shiny/lustrous
  Activity 3.11
  Demonstration of shininess in metals 
  Procedure: 
  1. Hold a small piece of sodium metal using forceps.
  2. Place it on a hard surface and cut it into two parts
  3. Observe the cut surface. What do you observe?
  4. Look at the surface of aluminium sheets, how does it appear? 
  Study question:
  Explain what makes metal surfaces appear shiny/luster.
  Light is composed of very small packages of electromagnetic energy called photons. 
  We are able to see objects because light photons from the sun (or other light source) 
  reflect off of the atoms within the object and some of these reflected photons reach 
  the light sensors in our eyes and we can see the objects. 
  When photons of light hit the atoms within an object three things 
  can happen: 
  The photons can bounce back from the atoms in the object, can pass through an 
  object such as glass or can be stopped by the atoms within the object.
  Objects that reflect many photons into our eyes make the objects appear shiny. 
  Objects that absorb photons and reflect less photons appear dull or even dark black 
  to our eyes.
  Did you know? Of all of the metals, aluminium and silver are the shiniest to our 
  eyes. Gold is also one of the more shiny metals. However, gold is not as shiny as silver 
  and aluminium. Mercury, a liquid metal, is also shiny and special telescope mirrors 
  have been made of mercury.
  (e)Melting and boiling points
  Activity 3.12:
  1. Why do metals have variable melting points?
  2. Why do metals have high melting points compared to non-metals?
  Melting point is a measure of how easy it is to separate individual particles. In metals 
  it is a measure of how strong the electron cloud holds the positive ions. The ease of 
  separation of ions depends on the electron density and Ionic / Atomic size.
  
  
  3.3.3. Factors affecting the strength of metallic bonds.
  Activity 3.13
  1. Explain different strengths of metallic bonds in different metals?
  2. Compare the metallic strength of the following metals:
  (i) Sodium and magnesium
  (ii) Sodium and potassium
  The three main factors that affect the metallic bond are:
  • Number of protons/ Strength of nuclear attraction: The more protons 
  the stronger the force of attraction between the positive ions and the 
  delocalized electrons
  • Number of delocalized electrons per atom: The more delocalized electrons 
  the stronger the force of attraction between the positive ions and the 
  delocalized electrons
  • Size of atom: The smaller the atom, the stronger the force of attraction 
  between the positive ions and the delocalized electrons and vice-versa, the 
  larger the atom, the weaker the force of attraction between the positive 
  ions and the delocalised electrons
  
  The strength increases across a period from left to right because:
  The atoms have more protons. There are more delocalized electrons per atom.
  Electrons are added to the same energy level. Group1 elements have 1 electron 
  in their outer shells and so contribute 1 electron to the sea of electrons, Group 2 
  elements contribute 2 electrons per atom, and Group 3 elements contribute 3 
  electrons per atom
  If the atoms/ions are smaller; there is therefore a greater force of attraction between 
  the positive ions and the delocalized electrons.
  In group 1 elements, the melting and boiling points decrease as the size increases 
  hence attraction between the delocalized electrons and metal cations decreases 
  

  down the group as shown table 3.6

  

  Checking Up 3.6
  1. Look at the table below, which shows the thermal conductivity of a number of 
  

  different materials, and then answer the questions that follow:

  

  The higher the number in the second column, the better the material is at 
  conducting heat (i.e. it is a good thermal conductor). Remember that a material 
  that conducts heat efficiently will also lose heat more quickly than an insulating 
  material. Use this information to answer the following questions: 
  1. Name two materials that are good thermal conductors.
  2. Name two materials that are good insulators. 
  3. Explain why: 
  a. cooler boxes are often made of polystyrene
  b. Homes that are made from wood need less internal heating during the 
  cold months.
  c. Igloos (homes made from snow) are so good at maintaining warm 
  temperatures, even in freezing conditions.
  d. Houses covered by iron sheets and houses covered by tiles can be 
  compared in their capacity of keeping the interior of the house hot of 
  fresh during a sunny and hot day.
  4. Magnesium has a higher melting and boiling point than sodium. This can 
  be explained in terms of the electronic structures, the packing, and the 
  atomic radii of the two elements. 
  a. Explain why each of these three things causes the magnesium melting 
  and boiling points to be higher. 
  

  b. Explain why metals are good conductors of electricity. 
  c. Explain why metals are also good conductors of heat. 
  3. Pure metals are usually malleable and ductile.
  a. Explain what those two words mean. 
   If a metal is subjected to a small stress, it will return to its original shape when 
  the stress is removed. However, when it is subjected to a larger stress, it may 
  change shape permanently. Explain, with the help of simple diagrams why 
  there is a different result depending on the size of the stress.
  When a piece of metal is worked by a blacksmith, it is heated to a high 
  temperature in a furnace to make it easier to shape. After working it with a 
  hammer, it needs to be re-heated because it becomes too difficult to work. 
  Explain what is going on in terms of the structure of the metal. 
  

  Why is brass harder than either of its component metals, copper and zinc?

  3.4. End Unit Assessment
  1. Choose from a list of words and fill in the missing words in the text below:
  List of words: 
  Conduct electricity, electrodes, electrolysis, electrostatic attraction, free electrons, 
  good conductivity, great malleability, high density,      high melting points, 
  ionic bond,metal, negative ion, non-metal, positive ion,regular crystal shape 
  and  attractive forces.
  Text:
  Metals have a layered structure of  ................... in fixed positions but between 
  them are oppositely charged ............... that can move around at random between 
  the metal atoms. There is a strong ...................................... between these oppositely 
  charged particles which gives them ..........................The strong forces also give 
  a ....................... making the average ........... heavier than an average ................ The 
  presence of ........................... in the structure keeps the bonding intact when metals 
  are bent or hammered giving them.......................................... Also, these ....................
  give metals ..............as regards heat and electricity.
  When electrons are transferred from (usually) ............  atom (e.g. sodium) 
  to.................  atom (e.g. chlorine) an ionic bond  is formed. Sodium loses an 
  electron to form a singly charged ........................... and chlorine gains an electron 
  to form a singly charged negative ion. In an ionic compound, the ionic bond 
  is the  electrostatic attraction  between the neighbouring positive ions and 
  

  negative ions. 

  The strong forces holding this giant ionic lattice together give 
  these ionic compounds............................ and...................................................................
  When ionic compounds are melted they are found to  ................  in a process 
  called  ...........using electrical contacts called............ In this process, move 
  to the negative electrode (cathode) and  metalsare released. At the same 
  time, ....................move to the positive electrode (anode) and ................ are formed. 
  Research from the internet or text books to find out other physical properties of 
  metals and ionic compounds that are not mentioned above.
  Answer these questions by choosing the best alternative represented by letters from A, B, C and D.
  1. Metals lose electrons from their lattice to become
  a. positive ions
  b. negative ions
  c. alkalis
  d. non- metals 
  2. Neither ions nor electrons are free to move in
  a. liquids
  b. metals
  c. ionic solids
  d. All of the above
  3. Attractive forces between metal ions and delocalized 
  electrons can be weakened or overcome by
  a. hammer
  b. high temperature
  c. water
  d. All of the above
  4. Metals are good conductors due to
  a. ionic lattices
  b. crystalline lumps
  c. mostly solids
  

  d. delocalized electrons

  5. Most atoms adopt one of three simple strategies to achieve a 
  filled shell. Which of the following is NOT one of these strategies?
  a. They accept electrons 
  b. They share electrons
  c. They give away electrons 
  d. They keep their own electrons
  6. Which of the following is NOT a type of chemical bond?
  a. Covalent
  b. Metallic
  c. Valence
   d. Ionic
  7. In metallic bonding...
  a . One atom takes the outer shell electrons from another atom.
  b. A couple of atoms share their electrons with each other.
  c. Some electrons are shared by all the atoms in the material. 
  d. Bonding takes place between positively charged areas of one atom with 
  a negatively charged area of another atom.
  8. Which of the following is NOT a characteristic of metals?
  A. Shiny /lustre
  B. Brittle/Shatters easily
  C. Conducts electricity
  D. Malleable
  9. When two or more metal elements are combined they form 
  an...
  a. bronze
  b. alloy
  c. Covalent bond
  d. Brass
  10. Sulphur is a solid non-metallic element at room temperature, so it is?
  a. A good conductor of heat
  b. A substance with a low melting point
  c. Easily bent into shape
  d. A good conductor of electricity
  11. Copper is a metallic element so it is likely to be a? 
  a. substance with a low boiling point
  b. poor conductor of electricity
  c. good conductor of heat
  d. substance with a low melting point
  12. Sodium chloride is a typical ionic compound formed by 
  combining a metal with a non-metal. Sodium chloride will? 
  a. have a low melting point
  b. consist of small NaCl molecules
  c. conduct electricity when dissolved in water
  d. not conduct electricity when molten
  13. Copper is a metallic element so it is likely to be a? 
  a. Substance with a shiny surface
  b. Poor conductor of electricity
  c. Poor conductor of heat
  d. Substance with a low melting point
  14. When an ionic bond is formed between atoms of different 
  elements? 
  a. Protons are transferred
  b. Electrons are transferred
  c. Protons are shared
  d. Electrons are shared
  15. Sodium chloride has a high melting point because it has: 
  a. Many ions strongly attracted together
  b. Strong covalent double bonds
  c. A giant covalent 3-dimentional structure
  d. Molecules packed tightly together
  16. Which substance is likely to have a giant ionic structure: 
  a. Melts at 1400o
  C, insoluble in water, good conductor of electricity either when 
  solid or molten
  b. Melts at 2800o
  C, insoluble in water, non-conductor of electricity when 
  molten or solid
  c. Melts at 17o
  C, insoluble in water, non-conductor of electricity either when 
  solid or molten 
  d. Melts at 2600o
  C, dissolves in water, non-conductor of electricity when 
  solid,undergoes electrolysis in aqueous solution
  17. Sodium chloride conducts electricity when:
  a. Solid or molten
  b. Solid or in solution 
  c. Molten or in solution
  d. Non of the above
  18. The structure of magnesium oxide is a
  a. Giant covalent lattice
  b. Giant ionic lattice 
  c. Simple ionic lattice 
  d. All the above
  19. What is the formula for magnesium chloride (contains Mg2+ and Cl?
   ions)?
  a. MgCl
  b. Mg22Cl
  c. MgCl2
  d. MgCl
  20. Why does sodium chloride have a lower melting point than magnesium 
  chloride?
  a. Its positive ions are smaller and have a smaller charge
  b. Its positive ions are larger but have a smaller charge
  c. Its positive ions are smaller but have a larger charge
  d. All the above
  

  21. Explain the conductivity of sodium chloride

  
  

  a. It conducts electricity when molten because it contains free electrons
  b. It conducts electricity when molten because sodium has metallic 
  bonding
  c. It conducts electricity when molten because its ions are free to move.
  d. None of the above
  Short and long answer questions
  22.(a) Explain why the lattice dissociation enthalpy of NaBr is a bit less than 
  that of NaCl.
  (b) Explain why the lattice dissociation enthalpy of MgO is about 5 times 
  greater than that of
  NaCl
  23.a) The table (using figures for lattice energies from gives experimental and 
  theoretical values for the silver halides.(The values are listed as lattice dissociation 
  

  energies.) compare the values and give a detailed explanation. 

  
  
  
  

  
  
  
  
  

  
  
  
  

  
  

  
  

  
  

• UNIT 4:COVALENT BOND AND MOLECULAR STRUCTURE

  UNIT 4: COVALENT BOND AND MOLECULAR 
  STRUCTURE
  Key Unit Competence
  Demonstrate how the nature of the bonding is related to the properties of covalent 
  compounds and molecular structures.
  Learning objectives
  By the end of this unit, students should be able to:
  • Define octet rule as applied to covalent compounds. 
  • Explain the formation of covalent bonds and 
   the properties of 
  covalent compounds. 
  • Describe how the properties of covalent compounds depend on their 
  bonding.
  • Explain the rules of writing proper Lewis structures
  • Draw different Lewis structures
  • State the difference between Lewis structures from other structures.
  • Apply octet rule to draw Lewis structures of different compounds. 
  • Make the structures of molecules using models. 
  • Write the structures of some compounds that do not obey octet rule.
  • Explain the formation of dative covalent bonds in different molecules.
  • Compare the formation of dative covalent to normal covalent bonding.
  • Describe the concept of valence bond theory.
  • Relate the shapes of molecules to the type of hybridization.
  • Differentiate sigma from pi bonds in terms of orbital overlap and formation.
  • Explain the VSEPR theory.
  • Apply the VSEPR theory to predict the shapes of different molecules/ions.
  • Predict whether the bonding between specified elements will be primarily 
  covalent or ionic.
  • Relate the structure of simple and giant molecular covalent compounds to 
  their properties.
  • Describe simple and giant covalent molecular structures.
  • Describe the origin of inter-molecular forces. 
  • Describe the effect of inter and intra molecular forces on the physical 
  properties of certain molecules. 
  • Describe the effect of hydrogen bonding in the biological molecules.
  • Relate the physical properties to type of inter and intra molecular forces in 
  molecules. 
  

  • Compare inter and intra molecular forces of attraction in different molecules. 

  

  In UNIT 3, you have learnt that atoms have different ways of combination to achieve 
  the stable octet electronic structure; two of those ways of combination led to 
  the formation of ionic bond and metallic bond. But what happens where the two 
  combining atoms need electrons to complete the octet structure and no one is 
  willing to donate electrons? For example the combination of 2 hydrogen atoms or 
  the combination of 2 chlorine atoms?
  When this happens, the combining atoms share a pair of electrons where each 
  atom brings or contributes one electron. In other words there is an overlapping of 
  two orbitals, one orbital from one atom, each orbital containing one electron (see 
  Fig.4.1): this bond is called “Covalent bond”. The attraction between the bonding 
  pair of electrons and the two nuclei holds the two atoms together.
  The covalent bond is a bond formed when atoms share a pair of electrons to complete 
  the octet. Similarly, people need each other irrespective of their race, economic, 
  political and social status for the success of human race. Some compounds that 
  exist in nature such as hemoglobin in our blood, chlorophyll in plants, paracetamol,
  responsible for transport of oxygen, green color in plants and as pain killer respectively 
  are made of the covalent bond. The covalent bonds mostly occur between nonmetals or between two of the same (or similar) elements.Two atoms with similar 
  electronegativity do not exchange an electron from their outermost shell; the atoms 
  instead share electrons so that their valence electron shell is filled.
  In general, covalent bonding occurs when atoms share electrons (Lewis model), 
  concentrating electron density between nuclei. The build-up of electron density 
  between two nuclei occurs when a valence atomic orbital of one atom combines 
  with that of another atom (Valence bond theory).In Valence bond theory, the bonds 
  are considered to form from the overlap of two atomic orbitals on different atoms, 
  each orbital containing a single electron.
  The orbitals share a region of space, i.e. they overlap. The overlap of orbitals allows 
  two electrons of opposite spin to share the common space between the nuclei, 
  

  forming a covalent bond.

  

  These two electrons are attracted to the positive charge of both the hydrogen 
  nuclei, with the result that they serve as a sort of ‘chemical glue’ holding the two 
  nuclei together. 
  The figure (Figure 4.1) shows the distance between the two nuclei. If the two nuclei 
  are far apart, their respective 1s-orbitals cannot overlap and no covalent bond is 
  formed. As they move closer each other, the orbital overlapping begins to occur, and 
  a bond starts to form. 
  The examples below represent different atoms overlapping in order to form covalent 
  

  bonds.

  

  

  

  4.1.1 Properties of covalent molecules
  Covalent molecules are chemical compounds in which atoms are all bonded together 
  through covalent bonds. The covalent compounds possess different properties and 
  some are emphasized below.
  • Covalent compounds exist as individual molecules, held together by weak 
  van der Waals forces.
  • Due to the weak van der Waals forces that hold molecules together, covalent 
  compounds have low melting and boiling points; because the weak forces 
  between molecules can be broken easily to separate the molecules. That 
  is why covalent compounds can be solid, liquid and gaseous at room 
  temperature. 
  • Covalent compounds do not display the electrical conductivity either in 
  pure form or when dissolved in water. This can be explained by the fact 
  that the covalent compounds do not dissociate into ions when dissolves in 
  water.
  • Generally non-polar covalent compounds do not dissolve in water; but 
  many polar covalent compounds are soluble in water( a polar solvent)
  • Non-polar covalent compounds are soluble in organic solvents (themselves 
  non-polar covalent). 
  The two statements above are at the origin of the say by chemists: “Like dissolves 
  

  like”

  

  

  bonding is usually placed at the center. The number of bonding sites is determined 
  by considering the number of valence electrons and the ability of an atom to expand 
  its octet. As you will progress in your study of chemistry, you will be able to recognise 
  

  that certain groups of atoms prefer to bond together in a certain way!

  

  

  electrons in the molecule or ion. In the case of a neutral molecule, this is nothing 
  more than the sum of the valence electrons on each atom. If the molecule carries 
  an electric charge, we add one electron for each negative charge or subtract an 
  electron for each positive charge.
  In Lewis structure, the least electronegative element is usually the central element, 
  except H that is never the central element, because it forms only one bond. 
  Another way of finding Lewis structure
  1. Calculate n (the number of valence (outer) shell electrons needed by all atoms in 
  the molecule or ion to achieve noble gas configurations for instance,
  NO3
  -
  , n=1× 8(for N atom) + 3×8 (for O atom) = 32 electrons.
  2. Calculate A, number of electrons available in the valence (outer) shells of all the 
  atoms. For negatively charged ions, add to this number the number of electrons 
  equal to the charge of the anions. For cations you subtract the number of electrons 
  equal to the charge on the cation. 
  For instance: NO3
  -
  , 
  A= 1×5(for N) +3×6 (for O atom) +1(for -1 charge) = 5+18+1=24 electrons.
  3. Calculate S, total number of electrons shared in the molecule or ion, using the 
  relationship
  S = n-A
  S= n-A= 32-24 =8 electrons shared (4pairs of electron shared)
  4. Place S electrons into the skeleton as shared pairs. Use double and triple bonds 
  only when necessary. Lewis formulas may be shown as either dot formula or dash 
  

  formulas.

  

  

  

  There are three general ways in which the octet rule doesn’t work:
  • Molecules with an odd number of electrons
  • Molecules in which an atom has less than an octet
  • Molecules in which an atom has more than an octet
  a. Odd number of electrons
  Consider the example of the Lewis structure for the molecule nitrous oxide (NO):
  Total electrons: 6+5=11
  Bonding structure:
  Octet on “outer” element is realized and on central atom only 3 electrons remain free 
  

  (11-8 = 3).

  

  

  

  

  

  4.3 Coordinate or dative covalent bonding and properties 

  

  4.3.1. Co-ordinate or dative covalent bonding and properties
  A dative covalent bond, or coordinate bond is another type of covalent bonding. In this 
  case, the shared electron pair(s) are completely provided by one of the participants 
  in the union, and not by contributions from the two of them. The contributors of the 
  shared electrons are either neutral molecules which contain lone pair(s) of electrons 
  on one of their atoms, or negatively charged groups with free pairs of electrons to 
  

  donate for sharing

  

  

  

  The solid copper (II) hydroxide which was initially formed reacts with the excess 
  ammonia (which acts as ligands) to form the water soluble tetra ammine copper (II) 
  

  complex as shown below. 

  

  

  4.4 Valence bond theory (VBT)

  
  

  
  

  
  

  As you notice, the density of bonding electrons is not on the inter-nuclei axis, it is 
  rather located outside the axis but surrounding it. This kind of covalent bond is called 
  “ Pi bond”, represented by the symbol “π”. Hence the double bond O=O is made of 
  two covelent bonds: a σ bond and a π bond.
  Due to the position of their electrons density in relation with the two nuclei, σ bond 
  participates in maintaining the two nuclei together more strongly than the π bond; 
  that is why σ bond is stronger than π bond. In addition, π bond cannot exist alone, it 
  exists only where there is a double or triple bond. Hence, in a double or triple bond, 
  there is one σ bond and one or two π bonds respectively. 
  Checking Up 4.4
  1. Describe the aspects and postulates of valence bond theory(VBT)
  2. Use VBT to explain the formation of single(sigma) and double (pi)bonds
  (a) Explanation of lateral overlap of atomic orbitals and
  

  (b) Explanation of head-to-head overlap of atomic orbitals

  4.5 Valence Shell Electron Pair Repulsion Theory (VSEPR) 

  theory

  
  

  
  

  nucleus. Hence they occupy more space. As a result, the lone pairs cause more 
  repulsion.
  The order of repulsion between different types of electron pairs is as follows: 
  Lone pair - Lone pair > Lone Pair - Bond pair > Bond pair - Bond pair
  The bond pairs are usually represented by a solid line, whereas the lone pairs are 
  represented by a lobe with two electrons.
  3) In VSEPR theory, the multiple bonds are treated as if they were single bonds. 
  The electron pairs in multiple bonds are treated collectively as a single super pair. 
  The repulsion caused by bonds increases with increase in the number of 
  bonded pairs between two atoms i.e., a triple bond causes more repulsion 
  than a double bond which in turn causes more repulsion than a single bond. 
  4) The shape of a molecule can be predicted from the number and type of 
  

  valence shell electron pairs around the central atom.

  
  

  The principle of the VSEPR is based on the idea that: the most stable structure of a 
  molecule is the one where the electron pairs are far away one from another in 
  order to minimize the repulsions between the pairs of electrons surrounding 
  the central atom.
  The VSEPR theory assumes that each atom in a molecule will achieve a geometry 
  that minimizes the repulsion between electrons in the valence shell of that atom. 
  The use of VSEPR involves the following steps:
  • Draw a Lewis structure for the ion or molecule in question.
  • The shape is based on the location of the nuclei in a molecule, so double 
  and triple bonds count as one shared pair when determining the shape of 
  the molecule
  • Locate the shared pairs and lone pairs on the central atom
  

  • Determine the shape based on the above considerations.

  

  

  4.6. Hybridisation and types of Hybridisation

  Activity 4.6

  1. Write the electronic configuration of carbon and hydrogen using s,p, d.. Notation.
  2. Use the electronic configurations above to identify the orbitals that contain 
  electrons used during the formation of methane. 
  3. Use the knowledge of overlap of atomic orbitals to indicate how orbitals overlap 
  in formation of hydrogen chloride, methane and beryllium chloride and predict 
  

  the shapes of the molecules.

  
  

  
  

  
  

  
  

  
  

  
  

  
  

  
  

  
  

  
  

  
  

  
  

  4.7 Polar covalent bonds
  Activity 4.7
  1. Can you define the term electronegativity?
  2. How is electronegativity related to polarity of the compound?
  3. How does the polarity of a given molecule affect its physical properties?
  4. Can you describe the general trends of electronegativity across and down 
  the groups in the periodic table?
  5. What is meant by the term dipole and Net dipole
  What happens if shairing of the bonding pair of electrons between the two atoms 
  forming the bond is not equal? For instance when two different non-metal elements 
  such as hydrogen and bromine combine?
  In this case, there is unequal sharing where the more electronegative element takes 
  a bigger share of the bonding pair of electrons (Fig. 4.7
  
  Figure 4.7: Polar covalent bond
  (www.bbc.co.uk/bitesize/higher/chemistry/energy/bsp/revision/1/
  In a polar covalent bond, binding pair of electrons is unequally shared between two 
  atoms. The power of an atom to attract the pair of electrons that constitutes the 
  bond in a molecule is called “electronegativity”.
  The ‘electronegativity’ can be used to determine whether a given bond is nonpolar covalent, polar covalent or ionic bond. The electronegativity increases from left 
  to right across a period and decreases as you go down a group
  The larger the electronegativity, the greater is the strength to attract a bonding pair 
  of electrons; and the larger the difference in electronegativites of the atoms, the 
  

  more polar the covalent bond between the two atoms.

  

  The following is the general thumb rule for predicting the type of bond based 
  upon the electronegativity differences:
  • If the electronegativities are equal and the difference in electronegativity 
  difference is less than 0.5, the bond is non-polar covalent.
  • If the difference in electronegativities between the two atoms is greater 
  than 0.4, but less than 2.0, the bond is polar covalent.
  • If the difference in electronegativities between the two atoms is 2.0, or 
  greater, the bond is ionic
  
  
  
   
  
  (ii) Poor electrical conductivity
  There are no charged particles (ions or electrons) delocalized throughout the 
  molecular crystal lattice to conduct electricity. They cannot conduct electricity in 
  either the solid or molten state.
  (iii) Solubility
  Simple structures tend to be quite insoluble in water, but this depends on how the 
  polarized molecule is. The more polar the molecules, the more water molecules will 
  be attracted to them (some may dissolve in water as a result of forming hydrogen 
  bonds within it). Molecular crystals tend to dissolve in non-polar solvents such as 
  alcohol.
  (iv) Soft and low density
  Van der Waals forces are weak and non-directional. The lattice is readily destroyed 
  and the crystals are soft and have low density.
  b. Giant covalent structures and their physical properties
  Sometimes covalently bonded structures can form giant networks, known as Giant 
  Covalent Structures. In these structures, each network of bonds connects all the 
  atoms to each other.
  These structures are usually very hard and have high melting and boiling points. 
  This is because of the strong covalent bonds holding each atom in place. In general, 
  Giant Covalent Structures cannot conduct electricity due to the fact that there are no 
  free charge carriers. One notable exception is Graphite. This is a structure composed 
  of ‘sheets’ of carbon atoms on top of each other. Electrons can move between the 
  sheets and carry the electricity. The main giant covalent molecular structures are the 
  two allotropes of carbon (diamond and graphite), and silica (silicon dioxide).
  (i) Diamond structure and the physical properties
  Diamond is a form of carbon in which each carbon atom is joined to four other 
  carbon atoms, forming a giant covalent structure with four single bonds. As a result, 
  diamond is very hard and has a high melting point. It does not conduct electricity. 
  Diamond is tetrahedral face-centered cubic as shown in the figure below
  
  Diamond has a very high melting point (almost 4000°C): the carbon-carbon covalent 
  bonds are very strong and have to be broken throughout the structure before 
  melting occurs.
  The compound is very hard due to the necessity to break very strong covalent bonds 
  operating in 3-dimensions.
  Diamond does not conduct electricity: All the electrons are held tightly between the 
  atoms, and are not able to move freely.
  The compound is insoluble in water and other organic solvents due to no possible 
  attractions which could occur between solvent molecules and carbon atoms which 
  could outweigh the attractions between the covalently bound carbon atoms.
  (ii) Graphite and the physical properties
  Graphite is another form of carbon in which the carbon atoms form layers. These 
  layers can slide over each other and graphite is much softer than diamond. Each 
  carbon atom in a layer is joined to only three other carbon atoms in hexagonal rings 
  

  as shown in the figure below

  

  

  Silicon dioxide exhibits some physical properties such as:
  • It has a high melting point (around 1700°C) which varies depending on 
  what the particular structure is (remember that the structure given is only 
  one of three possible structures).The silicon-oxygen covalent bonds are 
  

  very strong and have to be broken throughout the structure before the 
  melting occurs
  • Silicon dioxide is hard due to the need to break the very strong covalent 
  bonds.
  • Silicon dioxide is not displaying the property of electrical conductivity
  because all the electrons are held tightly between the atoms, and are not 
  able to move freely. No any delocalized electrons are observed.
  • It is insoluble in water and organic solvents because no possible attractions 
  occur between solvent molecules and the silicon or oxygen atoms which 
  could overcome the covalent bonds in the giant structure
  
  4.
  a) Draw a diagram to show the structure of silicon dioxide.
  b) Explain why silicon dioxide 
  (i) is hard;
  (ii) has a high melting point;
  (iii) Doesn’t conduct electricity;
  

  (iv) is insoluble in water and other solvents.

  4.9. Intermolecular Forces

  Activity 4.9
  1. Make a research and describe why:
  i) Ice floats over water and the bottle full of water breaks on cooling(freezing)
  ii) Water is a liquid at room temperature while Hydrogen sulfide is a gas
  2. Trichloromethane (ii) ethanol (iii) aluminium fluoride. Arrange these compounds 
  in order of increasing boiling points.
  Intermolecular forces are electrostatic forces which may arise from the interaction 
  between partial positively and negatively charged particles. Intermolecular forces 
  exist between two molecules while intramolecular forces hold atoms of a molecule 
  together in a molecule (Figure 4.11).
  Intermolecular forces are much weaker than the intramolecular forces of attraction 
  but are important because they determine the physical properties of molecules 
  such as their boiling point, melting point, density, and enthalpies of fusion and 
  

  vaporization.

  

  Intramolecular forces hold the atoms in the molecule together; they are called 
  chemical bonds. Intermolecular forces hold covalent molecules together and are 
  responsible of a certain number of properties of the substance such as the melting 
  and boiling temperatures of covalent substances. They can be grouped in a category 
  of forces called van der Waals forces. There are three main kinds of intermolecular 
  interactions such as London dispersion forces, dipole-dipole interactions and 
  

  hydrogen bonding later in the unit

  

  dispersion forces than chlorine, contributing to increasing the boiling point of 
  bromine, 59 o
  C, compared to chlorine, –35o
  C. Those London forces are very weak 
  for non-polar covalent compounds; hence breaking them does not require much 
  energy, which explains why non-polar covalent compounds such as methane and 
  nitrogen which only have London dispersion forces of attraction between the 
  

  molecules have very low melting and boiling points.

  

  4.9.3. Hydrogen bond
  For a hydrogen bond to be possible, there are necessary conditions:
  

  • The first condition is that the molecule contains one group where hydrogen is

  

  

  

  Hydrogen bonds in DNA

  (https://www.easynotescards.com/notecard_set/59549)

  Hydrogen bonding in ice
  Each water molecule is hydrogen-bonded to 4 others in a tetrahedral formation. The 
  ice molecule has a “diamond-like” structure. When liquid water freezes, the hydrogen 
  bonds become more rigid, and the volume becomes larger than the liquid because 
  of empty space generated by the rigidity of solid water. This explains why ice floats 
  

  over liquid water because its density is lower than the density of liquid water

  

  

  4.10. End unit assessment
  PART 1: MULTIPLE CHOICE QUESTIONS
  1. In the periodic table, electronegativity generally decreases: A From right to left 
  in a period, B Upwards in a group; C From left to right in a period.
  2. Which structure would sulphur, S8
  , have?
  A simple covalent molecules B simple covalent lattice C giant covalent lattice D
  giant ionic lattice
  3. Which statement(s) is/are true?
   1) Water has hydrogen bonds which increase the boiling point.
  2) Water as a solid is denser than as a liquid.
  3) Water has bond angles of 180o
  .
  A 1, 2 and 3
  B 1 and 2 
  C 2 and 3 
  

  D only 1

  

  a) The geometry of a molecule is determined by the number of electron groups 
  on the central atom.
  b) The geometry of the electron groups is determined by minimizing repulsions 
  between them. 
  c) A lone pair, a single bond, a double bond, a triple bond and a single electron - 
  each of these is counted as a single electron group. 
  d) Bond angles may depart from the idealized angles because lone pairs of 
  electrons take up less space than bond pairs. 
  e) The number of electron groups can be determined from the Lewis structure of 
  

  the molecule

  

  PART 2: Filling in questions
  16. Use the words listed below to fill in the correct appropriate word(s) in the 
  spaces below in the text.  
  Bigger, covalent bond, diamond, free electrons, halogens, hard crystals, 
  high electrical conductivity, high melting points, increase, intermolecular 
  forces, low electrical conductivity, low melting points, non-metals, 
  sharing, soft crystals, strong, strong bond, weak, weak force. 
  A ………………… is formed by two atoms…………….. one or more pairs of 
  electrons to make a …………………..between the two atoms in a molecule.
   However, between small molecules, only a ………………holds them 
  together in the bulk liquid or solid. This results in small covalent molecules 
  

  having ……………………… and ………………………………..if solid. Small 
  covalent molecules have no …………………. and so have a
  ……………………………………
  The Group 7 ………………………. collectively known as the…………………..
  form diatomic molecules of two atoms. The …………………………..between 
  the molecules are ……………. giving them relatively low melting points and 
  boiling points. This also explains why they are gases, liquids or solids with 
  ……………………. As you go down Group 7 the melting boiling points and 
  boiling points …………………..because the molecules get ………….. and the 
  intermolecular forces ………………….
  In giant covalent structures the forces between all the atoms are 
  ……………..forming ………………… like diamond or silica . In the 
  atomic giant structure metals there are free electrons which allow 
  …………………………………………….
  PART3: 
  17. Fill in the table by putting a check mark in the compare-and-contrast matrix 
  

  under the column(s) that each physical attribute describes.

  

  

  

  

  

  

  a) A 3-D, repeating pattern of + and – ions, formed by ionic compound
  b)Tendency for an atom to attract the bonding pair electrons when chemically 
  bonded to another atom
  c) A sharing of a pair of electrons
  d) Atoms will gain or lose enough electrons in order to become isoelectronic with 
  a noble gas
  d) A transfer of electrons from one atom to another
  e) A chemical formula that is arranged in the smallest whole number ratio
  f) The term that means dissolved in water
  g) A chemical formula that describes the makeup of a single molecule
  h) The shape (geometry) that is an exception to the octet rule
  i) A bond where electrons are shared unequally between atoms
  j) One of the shapes (geometries) that is polar
  k) A bond where electrons are shared equally between atoms
  
  

  Across

  11. A covalent bond between atoms in which the electrons are shared unequally
  12. A covalent bond in which the electrons are shared equally by the two atoms
  14. Intermolecular forces resulting from the attraction of oppositely charged 
  regions of polar molecules
  15. A bond formed when two atoms share a pair of electrons
  16. The two weakest intermolecular attractions - dispersion interactions and 
  dipole forces
  18. A covalent bond in which one atom contributes both bonding electrons
  20. A chemical formula that shows the arrangement of atoms in a molecule or 
  polyatomic ion
  21. A covalent bond in which three pairs of electrons are shared by two atoms
  22. A bond in which two atoms share two pairs of electrons
  23. A compound that is composed of molecules
  26. A molecule consisting of two atoms
  

  28. One of the two or more equally valid electron dot structures of a molecule or 
  polyatomic ion
  31. valence-shell electron-pair repulsion theory; because electron pairs repel, 
  molecules adjust their shapes so that valence electron pairs are as far apart as 
  possible
  Down
  1. An orbital that applies to the entire molecule
  2. A bond angle of 109.5 degrees that results when a central atom forms four 
  bonds directed toward the center of a regular tetrahedron
  3. The mixing of several atomic orbitals to form the same total number of 
  equivalent hybrid orbitals
  4. A tightly bound group of atoms that behaves as a unit and has a positive or 
  negative charge
  5. A pair of valence electrons that is not shared between atoms
  6. A molecule in which one side of the molecule is slightly negative and the 
  opposite side is slightly positive.
  7.a covalent bond in which the bonding electrons are most likely to be found in 
  sausage-shaped regions above and below the bond axis of the bonded atoms
  8. A covalent bond between atoms in which the electrons are shared unequally
  9. A neutral group of atoms joined together by covalent bonds
  10. A molecule that has two poles, or regions, with opposite charges
  12. The energy required to break the bond between two covalently bonded atoms
  17. A chemical formula of a molecular compound that shows the kinds and 
  numbers of atoms present in a molecule of a compound
  19. Attractions between molecules caused by the electron motion on one 
  molecule affecting the electron motion on the other through electrical forces
  24. Attractive forces in which hydrogen covalently bonded to a very electronegative 
  atom is also weakly bonded to an unshared electron pair of another electronegative 
  atom
  25. A solid in which all of the atoms are covalently bonded to each other
  27. A molecular orbital that can be occupied by two electrons of a covalent bond
  29. A bond formed by the sharing of electrons between atoms
  30. A bond formed when two atomic orbitals combine and form a molecular 
  orbital
  that is symmetrical around the axis connecting the two atomic nuclei
  
  
  

  
  
  

  
  
   
  
  
  
  

  

• UNIT 5:VARIATION IN TRENDS OF THE PHYSICAL PROPERTIES

  UNIT 5: VARIATION IN TRENDS OF THE PHYSICAL 
  PROPERTIES
  Key unit competence
  Use atomic structure and electronic configuration to explain the trends in the 
  physical properties of elements.
  Learning objectives
  By the end of this unit, students should be able to:
  • Outline the historical back ground of the Periodic Table. 
  • Explain the trends in the physical properties of the elements across a period 
  and down a group. 
  • Classify the elements into respective groups and periods using electronic 
  configuration. 
  • Relate trends in physical properties of the elements to their electronic 
  configuration. 
  

  • Classify the elements into blocks (s, p, d, f-block). 

  Introductory activity
  1. Explain how elements can be classified into a periodic table?
  2. Explain on which basis elements can be classified?
  3. How many groups and periods comprises a modern periodic table?
  4. How does electronic configuration of elements influence the structureof 
  modern periodic table?
  5. Discuss the basis of the location of the elements in the periodic table.
  5.1. Historical Background of the Periodic Table
  Activity 5.1
  Who is the father of the periodic table? Explain your answers.
  Differentiate the laws of triads and octaves
  During the nineteenth century, many scientists contributed to the development of 
  the periodic table. In the beginning, a necessary prerequisite to the construction of 
  the periodic table was the discovery of the individual elements. Although elements 
  such as gold, silver, tin, copper, lead and mercury have been known since antiquity, 
  the first scientific discovery of an element occurred in 1649 when Hennig Brand 
  discovered phosphorous. The periodic table of elements is a chart created in order 
  to help to organize the elements that had been discovered at that time. By 1869, 
  a total of 63 elements had been discovered. As the number of known elements 
  grew, scientists began to recognize patterns in properties and began to develop 
  classification schemes.
  Some important dates help us to understand more about how the periodic table has 
  been created. 
  • In 1669, Hennig Brand a German merchant and amateur alchemist invented 
  the Philosopher’s Stone; an object that supposedly could turn metals into 
  pure gold. He heated residues from boiled urine, and a liquid dropped out 
  and burst into flames. He also discovered phosphorus.
  • In 1680 Robert Boyle also discovered phosphorus without knowing about 
  Henning Brand’ discovery.
  • In 1809, curiously 47 elements were discovered and named, and scientists 
  began to design their atomic structures based on their characteristics.
  • In  1869, Dimitri Mendeleev  based on John Newlands’ ideas started 
  the development of elements organized into the  periodic table. The 
  arrangement of chemical elements were done by using atomic mass as the 
  key characteristic to decide where each element belonged in his table. The 
  elements were arranged in rows and columns. He predicted the discovery 
  of other elements, and left spaces open in his periodic table for them. At the 
  same time, Lothar Meyer published his own periodic table with elements 
  organized by increasing atomic mass.
  • In 1886, French physicist Antoine Becquerel first discovered radioactivity.
  During the same period of 1886, Ernest Rutherford named three types of 
  radiation; alpha, beta and gamma rays. 
  • In 1886, Marie and Pierre Curie started working on the radioactivity and 
  they discovered radium and polonium. They discovered that beta particles 
  were negatively charged.
  • In 1895, Lord Rayleigh discovered a new gaseous element named argon 
  which proved to be chemically inert. This element did not fit any of the 
  known periodic groups.
  • In 1898, William Ramsay suggested that argon be placed into the periodic 
  table between chlorine and potassium in a family with helium, despite the 
  fact that argon’s atomic weight was greater than that of potassium. This 
  group was termed the “zero” group due to the zero valency of the elements. 
  Ramsey accurately predicted the future discovery and properties neon.
  • In 1913, Henry Moseley worked on X-rays and determined the actual 
  

  nuclear charge (atomic number) of the elements. He has rearranged the 
  elements in order of increasing atomic number
  • In 1897 English physicist J. J. Thomson discovered small negatively charged 
  particles in an atom and named them as electrons;John Sealy Townsend
  and Robert A. Millikan investigated the electrons and determined their 
  exact charge and mass.
  • In 1900, Antoine Becquerel discovered that electrons and beta particles
  as identified by the Curies are the same thing.
  • In 1903, Ernest Rutherford  proclaimed that radioactivity is initiated by 
  the atoms which are broken down.
  • In 1911, Ernest Rutherford and Hans Geiger discovered that electrons are 
  moving around the nucleus of an atom.
  • In 1913, Niels Bohr suggested that electrons move around a nucleus in 
  discreete energy levels called orbits. He observed also that light is emitted 
  or absorbed when electrons transit from one orbit to another.
  • In  1914, Rutherford identified protons in the atomic nucleus. He also 
  transformed a nitrogen atom into an oxygen atom for the first time. 
  English physicist Henry Moseley provided atomic numbers, based on the 
  number of electrons in an atom, rather than based on atomic mass.
  • In  1932 James Chadwick discovered neutrons, and isotopes were 
  identified. This was the complete basis for the periodic table. In that same 
  year Englishman Cockroft and the Irishman Walton first split an atom by 
  bombarding lithium in a particle accelerator, changing it to two helium 
  nuclei.The last major changes to the periodic table give rise from Glenn 
  Seaborg’s work in the middle of the 20th Century. In 1940, he discovered 
  plutonium and all the transuranic elements from 94 to 102. 
  • In 1944, Glenn T. Seaborg discovered 10 new elements and moved out 14 
  elements of the main body of the periodic table to their current location 
  below the lanthanide series. These elements were known as Actinides series.
  • In 1951, Seaborg was awarded the Nobel Prize in chemistry for his work. 
  Element 106 has been named seaborgium (Sg) in his honor.
  

  • Presently, 118 elements are in the modern Periodic Table. 

  5.1.1. Law of Triads
  In 1817 Johann Dobereiner noticed that the atomic weight of strontium fell midway 
  between the weights of calcium and barium, elements possessing similar chemical 
  properties. 
  In 1829, after discovering the halogen triad (three) composed of chlorine, bromine, 
  and iodine and the alkali metal triad of lithium, sodium and potassium he proposed 
  that nature contained triads of elements the middle element had properties that 
  were an average of the other two members when ordered by the atomic weight (the 
  Law of Triads).
  Between 1829 and 1858 a number of scientists (Jean Baptiste Dumas, Leopold 
  Gmelin, Ernst Lenssen, Max von Pettenkofer, and J.P. Cooke) found that these types 
  of chemical relationships extended beyond the triad. During this time fluorine was 
  added to the halogen group; oxygen, sulfur, selenium and tellurium were grouped 
  into a family while nitrogen, phosphorus, arsenic, antimony, and bismuth were 
  classified as another. Unfortunately, research in this area was hindered by the fact 
  that accurate values were not always available. 
  5.1.2. Law of Octaves
  In 1863, John Newlands, an English chemist suggested that elements be arranged in 
  “octaves”. He wrote a paper in which he classified the 56 established elements into 
  11 groups based on similar physical properties, noting that many pairs of similar 
  elements existed which differed by some multiple of eight in atomic weight. This 
  law stated that any given element will exhibit analogous behavior to the eighth 
  element following it in the table. However, his law of octaves failed beyond the 
  element calcium.
  Although Dimitri Mendeleev is often considered the “father” of the periodic 
  table, however the work of many scientists contributed to its present form. The 
  representation of a modern Periodic Table of Elements is shown below.
  
  Checking up 5.1
  The periodic table is an important tool used in chemistry:
  

  1.Explain why the elements are classified in groups and periods of the periodic 
  

  table
  

  2. Chose one element of Group 1 and one of group 17 and make their electronic 
  configurations using orbitals.
  5.2. Comparison of Mendeleev’s Table and Modern Periodic 
  Table
  Activity 5.2
  1. Discuss the similarities and differences of Mendeleev’s table and modern 
  periodic Table.
  2. How were the positions of cobalt and nickel resolved in the modern periodic 
  table?
  The periodic table is the arrangement of chemical elements according to their 
  chemical and physical properties. The modern periodic table was created after 
  a series of different versions of the periodic table. The Russian Chemist/Professor 
  Dmitri Mendeleev was the first to come up with a structure for the periodic table with 
  columns and rows. This feature is the main building block for the modern periodic 
  table as well. The columns in the periodic table are called groups, and they group 
  together elements with similar properties. The rows in the periodic table are called 
  periods, and they represent sets of elements that get repeated due the possession of 
  similar properties. The main difference between Mendeleev and Modern Periodic 
  

  Table are shown in the Table below (Table 5.1).

  Table 5.1. Differences between Mendeleev’s table and the modern Periodic 

  Table

  
  

  
  

  Checking up 5.2

  1. The periodic table is an arrangement of elements based on their properties.
  Explain the gaps found in the Mendeleev periodic table compared to the modern 
  one?
  2. How many elements does the modern periodic table contain?
  3. Look at the modern periodic table and write down four things it tells you.
  5.3. Location of Elements in the Periodic Table Based On the 
  Electronic Configuration
  Activity 5.3
  1. Based on knowledge gained in the previous years:
  a. Represent the electronic configuration of the elements 25X and 11Y.
  b. Discuss the information given by the number of electrons in the last orbitals 
  of the above element about their position in the periodic table? 
  c. Explain the period and the group of the periodic table in which the above 
  elements are located.
  2. Is it possible to have an element with atomic number 1.5 between hydrogen 
  and helium?
  5.3.1. Major Divisions of the Periodic Table
  The periodic table is a tabular of the chemical elements organized on the basis of 
  their atomic numbers, electron configurations, and chemical properties.
  In the periodic table, the elements are organized by periods and groups. The period 
  relates to the principal energy level which is being filled by electrons. Elements 
  with the same number of valence electrons are put in the same group, such as the 
  halogens and the noble gases. The chemical properties of an atom relate directly to 
  the number of valence electrons, and the periodic table is a road map among those 
  properties such that chemical properties can be deduced by the position of an 
  element on the table. The electrons in the outermost or valence shell are especially 
  important because they participate in forming chemical bonds.
  Elements are presented in increasing atomic number. The main body of the table is 
  a 18 × 7 grid. There are four distinct rectangular areas or blocks such as s, p, d and 
  f blocks. The f-block is usually not included in the main table, but rather is floated 
  below, as an inline f-block would often make the table impractically wide. Using 
  periodic trends, the periodic table can help predict the properties of various elements 
  and the relations between properties. It therefore provides a useful framework for 
  analyzing chemical behavior and is widely used in chemistry and other sciences 
  (Petrucci et al., 2007).
  5.3.2. Location of elements in modern Periodic Table using examples
  

  In the periodic table, the elements are located based on groups and periods.

  

  

  

  

  

  receives the last electrons. The s-block has two groups of reactive metals: Group 1 
  and 2.
  p-block is composed of metals and nonmetals of Group 13 to 18.
  d-block is made of transition metals:Group 3 to Group 12, and f-block is made of 
  lanthanide and actinide series or inner transition metals.
  The division of elements into blocks is primarily based upon their electronic 
  configuration as shown in Figure 5.1. Two exceptions to this categorization can 
  be mentioned. Helium is placed in p-block although its valence electrons are in s 
  orbital because it has a completely filled valence shell (1s2
  ) and as a result, displays 
  properties representative of other noble gases. The other exception is hydrogen. It 
  has only one s-electron and hence can be placed in group 1 (alkali metals); but in 
  many modern Periodic Tables, hydrogen is left hanging above the Periodic Table and 
  doesn’t belong to any group. This is due to the particular properties of hydrogen:
  • Hydrogen is the smallest chemical element
  • Hydrogen is a gas while the other elements of group 1 are solids,
  • Hydrogen is not a metal whereas the other elements of group1 are metals, 
  • In some compounds where hydrogen combine with non-metals, it behaves 
  like a metal, e.g. in the polar molecule Hδ+Clδ-, hydrogen tends to lose an 
  electron,
  • When combined with very active metals, it behaves as a non-metal and 
  forms a negative ion H-
  , hydride ion; e.g. Na+H-
  (sodium hydride).
  Elements within the same group have the same number of electrons in their 
  valence (outermost) shells, and they have similar valence electron configurations. 
  They exhibit similar chemical properties. Elements within the same period have 
  different numbers of electrons in their valence shells and the number of electrons is 
  increasing from left to right. Therefore, elements in the same period are chemically 
  different, changing from metals to non-metals across the period from left to right.
  
  
  

   

   

  

  Checking up 5.5
  1. Among the common blocks, s, p, and d; which block has a tendency to form 
  complex compounds?
  2. Why d-block are called transition elements?
  3. Why f-block are called inner transition elements?
  5.6. Variation of Physical Properties down the Groups and 
  across the Periods
  Activity 5.6
  1. The elements in the periodic table display many trends which can be used to 
  predict their physical properties. Explain three of the factors that you think can 
  influence the physical properties of elements in the periodic table.
  2. Discuss the trends of the above factors across a period and down a group in the 
  periodic table.
  The elements in the periodic table are arranged in order of increasing atomic number. 
  All of these elements display several other trends and we can use the periodic table 
  to predict their physical properties. There are many noticeable patterns in the 
  physical and chemical properties of elements as we descend in a group or move 
  across a period in the Periodic Table. 
  Those trends can be observed in: ionization energy, electronegativity, electropositivity, 
  electron affinity,  melting and boiling point, density and metallic character and hereafter 
  are some factors which cause those trends.
  5.6.1. Atomic radius
  The atomic radius of an atom is defined as half the distance between the nuclei of 
  two atoms of the same element that are joined together by a single covalent bond.
  Atomic radius of elements decreases as we move from left to right in periodic table. 
  This is explained by the number of outer electrons and protons which increase while 
  there is no change in the energy level. The results increase the attracting forces 
  making the radius smaller.
  Increasing nuclear charge (more protons) pulls the electrons closer to the nucleus, 
  and the screening effect of inner electron shells will be the same for all members of 
  a given period. The combined effect of both factors results in the electrons being 
  pulled closer to the nucleus and a smaller radius. 
  On the other side, in the same group, as we go down, the atomic radius of elements
  increases. This is due to the energy level which increases when you move down in 
  group of the periodic table, the attraction of external electrons by nucleus decreases 
  and atomic radius increases.
  In general, atomic radii increase down a group because a new shell is added for 
  each successive member of a group, leading to a greater radius. Then an increased 
  screening effect of extra electron shells i.e. the nucleus has less of a pull on the outer 
  electrons. 
  5.6.2. Electronegativity
  Electronegativity is a measure of the tendency of an atom to attract to itself the 
  shared pair of electrons making a bond. The charge in the nucleus increases from 
  left to right across a period.The electronegativity of atoms is affected by both the 
  charge of the nucleus and the size of the atom. The higher its electronegativity, the 
  more an element attracts electrons. In general, the electronegativity of a non-metals 
  is greater than that of metals. Trends are observed in the period (Figure 5.3) or in a 
  group of the Periodic Table (Figure 5.4).
  • In a period, the electronegativity increases from left to right. This is 
  explained by the fact that as we go from left to right, there in an increase 
  of positive charge in the nucleus, since the number of protons increases; 
  but the electrons are being added to the same energy level. This results in 
  the reduction of the volume or radius of the atoms from left to right and 
  explains why attraction of external electrons by the nucleus increases from 
  left to right.
  • In a group, the electronegativity decreases from top to bottom. This is due 
  increase of energy levels down in a group, and thus there is an increased 
  distance between the valence electrons and the nucleus, or a greater atomic 
  radius.The positive charge of the nucleus is further away from the valence 
  

  electrons and the nucleus cannot attract efficiently external electrons.

  

  Note:
  • Since noble gases do not react or do not form chemical bonds, their 
  electronegativity cannot be determined.
  • For the transition metals, the electronegativity does not vary significantly 
  across the period and down a group. This is because their electronic structure 
  affects their ability to attract electrons easily like the other elements.
  • The lanthanides and actinides possess  more complicated chemistry 
  that does not generally follow any trend. Therefore, they do not have 
  electronegativity values.
  According to these two general trends, the most electronegative element is fluorine
  

  and Francium is the least (Figure 5.4 and Figure 5.5).

  

  The charge in the nucleus increases across a period. Greater is the number of protons, 
  greater is the attraction for bonding electrons.
  5.6.3. Ionization energy (I.E)
  Ionization energy: it is the amount of energy required to remove an electron from 
  a neutral gaseous atom. The lower this energy is, the more readily the atom loses 
  electron and becomes a cation. Therefore, the higher this energy is, the more 
  unlikely it is the atom to become a cation. We can distinguish, first, second, and third 
  ionization. Helium is the element with the highest ionization energy (Zumdahl and 
  Zumdahl, 2010). The noble gases possess very high ionization energies because of 
  their full valence shells compared to the elements of group 1 (Table 5.5).
  The table shows that generally the IE decreases down the Group, as the size of the 
  

  atoms increases down the Group.

  

  Ionisation energy of rare gases or any species with an octet electronic structure show 
  very high IE because the electron is being removed from a very stable electronic 
  structure.
  The ionization energy varies across a period and down a group.
  Across a period ionisation energies increase because the nuclear charge increases 
  (greater positive charge on the nucleus) and holds the outer electrons more strongly. 
  More energy needs to be supplied to remove the electron.
  Down a group ionisation energies decrease because the outer electrons are further 
  away from the nucleus. The screening effect of the inner electron shells reduces the 
  nuclear attraction for the outer electrons, despite the increased (positive) nuclear 
  charge.
  5.6.4. The melting points and boiling points
  Melting points and boiling points show some trends in groups and periods of the 
  Periodic Table.
  As you already know, the Periodic table can be subdivided into two main area or 
  regions: 
  • the left region where you find only metallic elements
  • the right region where you find both metallic and non-metallic elements; 
  

  all non-metallic elements are in the extreme right part of that region.
  The general trends of melting and boiling points depends on the regions:
  • in the left region, melting and boiling points generally decrease down the 
  groups due to the decrease of strength of the metallic bond down the 
  groups;
  • on the contrary, in the right region at the extreme right in groups 17 and 
  18, there is a general increase of melting and boiling points down the group 
  due to the increase of the molecular mass;
  • from left to the middle of the periodic table, there is an increasing of melting 
  and boiling points from left to right in a periode due the the increasing of 
  the strength of the metallic bond;
  • whereas from the middle of the periodic table, there is a decrease of melting 
  and boiling points from left to right due to the progressive increase of nonmetallic character where elements exist as simple molecules.
  The melting and boiling points vary in a regular way or pattern depending on their 
  position in the Periodic Table. In general the forces of attraction for elements on the 
  left of the table are strong metallic bonds; they require higher energy to be broken, 
  hence higher melting and boiling points.
  As we cross toward the right side of the periodic table, the non-metal character of 
  elements increases and elements, except few elements, form molecules that are 
  held together by weak intermolecular forces; hence their melting and boiling points 
  are generally low.
  For example going down in group 1, the melting point and boiling point of the 
  alkali metals decrease. This is due to the weakning of metallic bond down the group. 
  However, going down in group 17 of the halogens the melting point increases 
  meaning that there is an increase in the force of attraction between the molecules. 
  The illustrations below show the variation of melting and boiling point for some 
  

  elements of the periodic table (Figures 5.6 and 5.7).

  

  

  5.6.5. The density
  The density of a substance is its mass per unit volume, usually in g/cm3
  . The density 
  is a basic physical property of a homogeneous substance; it is an intensive property, 
  which means it depends only on the substance’s composition and does not vary 
  with size or amount.
  The trends in density of elements can be observed in groups and periods of the 
  periodic table. In general in any period of the table, the density first increases from 
  group 1 to a maximum in the centre of the period because the mass increases while 
  the size decreases, and then the density decreases again towards group 18 because 
  of the nature of bonds.
  Going down a group gives an overall increase in density because even though the 
  volume increases down the group, the mass increases more. 
  

  The variation of density with atomic number is shown in the Figure 5.8.

  

  5.6.6. Electrical and thermal conductivity
  The electrical conductivity is the ability of a substance to conduct an electric current.
  The electrical conductivity of elements increases from non-metals to metals. Metals 
  are good conductor of electricity. This is due to the presence of free electrons in 
  metallic lattice. The capacity of metals to conduct heat is called thermal conductivity 
  of metals. Electrical conductivity results from the transfer or mobility of electrons, 
  whereas the thermal conductivity in metal is due to heat transfer by free electrons 
  from one end of metal to another end.
  As we move across the period from the left to the right, the electrical conductivity 
  increases for the metals as the number of free electrons increases and then decreases 
  for the non-metals because they do not have free and mobile electron.
  1. Metallic character
  Metallic character refers to the level of reactivity of a metal. Metals tend to lose 
  electrons in chemical reactions, as indicated by their low ionization energies. Within 
  a compound, metal atoms have relatively low attraction for electrons, as indicated 
  by their low electronegativities.
  Metals are located in the left and lower three-quarters of the periodic table, and tend 
  to give electrons to nonmetals. Nonmetals are located in the upper right quarter of 
  the table, and tend to gain electrons from metal. Metalloids are located in the region 
  between the other two classes and have properties properties.
  • Metallic character is strongest for the elements in the leftmost part of the 
  periodic table and tends to decrease as we move to the right of any period.
  • Within any group of the representative elements, the metallic character 
  increases progressively going down.
  2. The electron affinity (E.A)
  The electron affinity is the ability of an isolated gaseous atom to accept an electron. 
  Unlike electronegativity, electron affinity is a quantitative measurement of the 
  energy change that occurs when an electron is added to a neutral gas atom. The more 
  negative the electron affinity value, the higher an atom’s affinity for electrons. In the 
  periodic table, the first electron affinities of elements are negative in general except 
  the group 18 and group 2 elements. The second electron affinities of all elements 
  are positive. This is because the negative ion creates a negative electric field. And 
  if now the other electrons enter the negative field, energy has to be applied to the 
  system to overcome the repulsion between the negative electric field and incoming 
  electron. 
  The more the electron affinity value is negative, the higher is the electron affinity of 
  an atom. Electron affinity decreases down a group of elements because each atom is 
  larger than the atom above it (refer to atomic radius trend).This means that an added 
  electron is further away from the atom’s nucleus compared with its position in the
  smaller atom. With a larger distance between the negatively-charged electron and 
  the positively-charged nucleus, the force of attraction is relatively weaker. Therefore, 
  electron affinity decreases down the group.
  Moving from left to right across a period, the electron affinity increases because 
  the electrons added to energy levels become closer to the nucleus and there is a 
  

  stronger attraction between the nucleus and electrons.

  Checking up 5.6

  1. The following table shows a part of a periodic table. Students have to answer 

  the following

  

  Fill in the blank space with the correct term based on the above table.

  The element with the least nuclear charge is …….and the one with the highest 
  nuclear charge …..Nuclear charge of S is ….than the nuclear charge of Se.
  As you go from Na to Cl along the period nuclear charge ….
  Effective nuclear charge….from B to Ga while it….from Na to Ar. Shielding or 
  screening effect ….. down the group but ……along the period from left to right.
  Atomic size of Li is…. than that of K. Element with the least atomic size is (10)……
  and the element with the highest atomic size is …….
  Atomic size of Ca is …. than atomic size of Be because number of …….increases 
  down the group.
  Atomic size …….. from K to Kr because electrons are filled on the same shell, the 
  …….continuously increase and attraction force increases.
  Analyze and complete the following concept map using: ionization energy, atomic 
  

  size, electron affinity, electronegativity and metallic character

  5.7. End unit assessment
  1. The following are coded groups/families of the representative elements of the 
  periodic table (first 4 periods, s, p blocks only). The groups are in number of 
  particular order. Use the hints below to identify the group and place of three 
  elements of each group in their correct location in the periodic table: AOU, BVW, 
  

  CKM, DLQ, ENT, FIJ, GPY, and HRS.

  

  Hints
  A has only one electron in p subshell
  B is more electronegative than V
  C has a larger atomic radius than both M and W
  D has electronic configuration ending in p5
  E is one of the most reactive metals
  F has a smaller ionization energy than J
  G has only 1 energy level with any electrons
  H has one more proton than O and is in the same period as O
  I is the largest alkaline earth metal
  J has one more proton than E
  K has electron configuration ending in p3
  L has more filled energy levels than D
  M is larger than K
  N has the largest radius in its family
  O is smaller than F but in the same energy level as F
  P is smaller than Y
  Q is the most reactive non-metal
  R has the highest electronegativity in its family
  T has the lowest density in its family
  U more easily loses electrons (think about ionization energy) than either A or O
  V has only 4 electrons in a p-subshell
  W has 3 completely filled energy levels
  Y has the lowest ionization energy in its family.
  2. Based on the variation of ionization energy in groups and periods, how should 
  you explain the variation of first and second ionization energy down a group and 
  across a period?
  3. Justify the following statements:
  a) The first ionization energy of nitrogen is higher than that of oxygen even though 
  nuclear charge of nitrogen is less compared to oxygen.
  b) Noble gases are having high ionization energies.
  4. Give reason
  a. Alkali metals (group 1 elements) are not found free in nature.
  b. Atomic radius of gallium is smaller than that of aluminium.(Z of Al = 13, Z of Ga 
  

  = 31

  
  
  

  
  

• UNIT 6:TRENDS IN CHEMICAL PROPERTIES OF GROUP 1 ELEMENTS AND THEIR COMPOUNDS

  UNIT 6: TRENDS IN CHEMICAL PROPERTIES OF GROUP 
  1 ELEMENTS AND THEIR COMPOUNDS
  Key unit competence: Compare and contrast the chemical properties of the Group 1 
  elements and their compounds in relation to their position in the Periodic Table.
  Learning objectives
  By the end of this unit, students should be able to:
  • Describe and explain the physical properties of Group 1 elements in terms 
  of metallic character and strength of metallic bond
  • Describe and explain the reactivity of Group 1 elements with oxygen, water 
  and halogens
  • State and explain the properties of Group 1 oxides and hydroxides
  • Explain the trends in the solubility of Group 1 compounds
  • State the uses of Group 1 elements and their compounds
  • Compare the reactivity of Group 1 elements
  • Interpret the trends in the thermal decomposition of Group 1 carbonates 
  and nitrates
  • Perform experiments to test the alkalinity of Group 1 hydroxides
  

  • Carry out flame test for the presence of Group 1 metal cations in solution.

  

  2. Identify elements of this group in question 1 above which are the most 
  and least reactive. 
  3. Give examples of compounds of some of the elements in the group in 
  question 1 above and suggest their uses. 
  6.1. Occurrence and physical properties of group 1 elements, 
  physical state, metallic character, physical appearance 
  and melting point
  Activity 6.1.a
  1. Study the following table of data and answer the questions that follow
  

  Data

  

  metals. For this reason, the study of hydrogen is presented separately from the other 
  members of group 1 elements.
  Francium is exceptionally rare. It is formed by the radioactive decay of heavier 
  elements.
  Because Francium is both rare and highly radioactive, few of its properties have been 
  determined and we are not going to talk much about it in this Unit.
  Physical Properties of Alkali Metals
  Activity 6.1.b
  In groups, learners make research in libraries / internet and discuss the physical 
  properties of group 1 elements and or explain the following statements
  a) Group 1 elements show weak metallic bonding
  b) Atomic radius of Na is smaller than the corresponding ionic radius of Na+
  c) The shining appearance of metals disappears after a certain period of time.
  Group 1 elements are grey metals, soft, and can be easily cut with a knife to expose 
  a shiny surface which turns dull on reaction with oxygen in air.
  1. They have low melting and boiling points. They show relatively weak metallic bonding as only one valence electron is attracted by the nucleus. Also 
  they have a big atomic radius and the attraction of nucleus toward the valence electron is weak.
  2. They are good conductors of heat and electricity 
  3. They have a low ionisation energies that decreases down the group
  4. They have low density compared to other metals and Li, Na and K are less 
  dense than water
  Group 1 metals color flames: When alkali metals are put in a flame they produce 
  
  
  
  Checking up 6.1
  1. Discuss how the ionization energies vary in function of atomic radius for 
  group 1 elements.
  2. Is there any relationship between the atomic radius and the melting 
  point in group 1 elements? Yes or No. Justify your answer 
  3. Why group 1 elements are said to be good conductors of electricity? 
  Illustrate your answer 
  4. The following table shows 3 unknown group 1 metals X, Y, Z and some of 
  their physical properties. Predict among the alkalis metal (Cs, Li, K), which 
  

  one should correspond to X,Y,Z. Justify your answer

  

  6.2. Reactivity of group 1 elements with oxygen, water and 
  halogens
  Activity 6.2 (a)
  Analyse the case study below and answer the related question:
  1. In groups learners discuss the following scenario and compare their 
  findings to the reactivity of a chemical element in terms of the variation 
  of atomic radius
  Suppose that a hen is walking in the garden with its chicks. Some of the chicks 
  are feeding themselves just near the hen (mother). Other chicks are feeding 
  themselves far away from the mother. Which ones of the two groups of chicks will 
  

  be an easy prey of a predator? Explain.

  

  Now extend your reasoning to the behavior of group 1 element and explain the 
  following statement
  Group 1 elements react by losing their single electron in outermost shell. Arrange 
  them in order to show which one loses easily the single electron and which one 
  loses electron with difficult.
  2. Consider a case of people who are warming around a fire; some are 
  near, others are a little bit far; who will feel the heat of the fire more than 
  

  the others?

  

  These activities show how the distance between two objects affects interactions 
  between them. The hen cannot protect the chicks that are far from her. In the other 
  case, people who are far from the fire feel less the heat of the fire. 
  This resembles the interactions between the nucleus and the valence electrons. The 
  attraction between the nucleus and the valence electron decreases with increasing 
  distance between the nucleus and the valence electrons.
  When the atomic radius and volume increase, the distance between the nucleus and 
  the valence electrons increases, the attraction between the nucleus and the valence 
  electrons decreases, and it becomes easier to remove the valence electrons.
  This explains the origin of properties of group 1 metals: they have low ionization 
  energy; lose easily the only one valence electron to form a mono-positive cation 
  (M+) with a rare gas electronic structure, and consequently are very active metals.
  Activity6.2 (b)
  a) Given the following elements of group 1(Na (Z=11), Li(Z=3), Cs(Z=55), K(Z=19) , 
  Rb (Z=37). Arrange them according to their increasing reactivity and justify why.
  b)Establish the electronic structure for the following species and explain how you 
  get it Na+,Na2+,Na-
  

  .Which one is stable and why?

  6.2.1. Reactions with Oxygen

  Activity 6.2.1

  Experiment: burning an alkali metal in air/oxygen

  
  

  Apparatuses: deflagrating spoon, Bunsen burner, glass beaker, filter paper,
  Chemicals: lithium, sodium, potassium, water, and red litmus paper
  Other requirements: knife, match box and petroleum gas
  Procedure:
  1. Cut a small piece of lithium and wrap it in piece of filter paper to remove the oil.
  2. Place it on to a deflagrating spoon and heat it in a non luminous flame.
  3. Observe what happens 
  4. When combustion is complete, dip the deflagrating spoon into a beaker of 100 
  ml filled up to 50 ml of water.
  5. Stir the water with the spoon and then drop a piece of litmus papers into the 
  solution in the beaker. Observe
  6. Repeat the experiment with sodium and potassium
  Task on the experiment:
  a) Write the equations of reactions that take place when each metal is burnt in air
  b) Name the product that was formed in each case.
  c) What are the color changes when the aqueous solutions above are tested with 
  litmus paper? Explain why?
  d) Write the equations of reaction between each product in (b) with water
  
  

  6.2.2. Reaction with water
  Activity 6.2.2 Experiment to investigate the reaction of alkali metals with 
  water
  Procedure:
  Cut a small piece of sodium metal and put it in water in wide beaker and observe.
  Test the obtained solution with red and blue litmus papers and observe.
  Questions 
  1. Why is sodium so easy to cut?
  2. What are the observations when sodium is placed in water in the beaker
  3. Explain the observation when tests with red and blue litmus papers are 
  performed on the resulting solution and write the equation of the reaction that 
  takes place to explain you answer
  
  6.2.3. Reaction with halogens
  Activity 6.2.3
  a) In terms of s, p, d, f orbitals write the electronic configuration of chlorine (Z=17), 
  bromine (Z=35) and iodine (Z=53)
  b) Deduce the valency of each element.
  c) Write molecular equations, complete and balance them, when
  Sodium reacts with bromine Potassium reacts with iodine
  

  Lithium reacts with chlorine

  

  Checking up 6.2
  1. An element J has 19 as atomic number while the element A has 35 as atomic 
  number.
  a) Write their electronic structures in term of s, p, d, f orbitals and deduce their 
  respective valencies.
  b) The element J is able to react with oxygen gas by forming two types of oxides
  i) Write the formula of the 2 oxides that can be formed between the element J and oxygen
  ii) Write the formula of the compound formed between J and A
  iii) What type of bond does exist between J and A. Justify your answer.
  c) Show how you would write the equation of reaction between J and water 
  supposing that J stands for the real symbol of the element.
  d) When the reaction stated in (c) takes place a colorless solution and a colorless 
  gas are formed.
  i) Which test would you use to identify each product of the reaction, by stating the 
  reagent and related observations?
  Oxides of Group 1 metals dissolve in water to give strong alkaline solutions; that is 
  why they are said to form basic oxides.
  6.3.2. Hydroxides
  As said earlier hydroxides are formed when the metals or metals oxides are dissolved 
  in water. In solid state, these hydroxides dissolve very easily in water and in alcohol. 
  They dissociate completely in water to form alkaline solution; hence they are strong 
  bases. The basic character of the hydroxides increases as we move down the group.
  

  Checking up 6.3
  

  1. Complete and balance the reactions when water is reacted with the following:
  a) Potassium metal b) Potassium oxide
  2. Lithium hydroxide decomposes on heating. White powder X and a colorless gas 
  Y is released and condenses in a colorless liquid.
  i) Write the chemical formula of X. ii) Propose a chemical test to identify Y 
  

  3. Explain why Group 1 metals form ionic compounds?

  6.4. The effect of heat on Group 1 carbonates and nitrates

  6.4.1. Heating the nitrates

  Activity 6.4.1

  Experiment: effect of heat on nitrates.
  In groups learners perform the following experiment, discuss and make 
  conclusions by explaining the observed phenomena, and write involved chemical 
  reactions.
  Apparatus: glass test tubes, pair of tongs, wooden splint/match stick, Bunsen 
  burner/heat source and spatula.
  Chemicals: Lithium nitrate, potassium nitrate
  

  Other requirements: match box

  Procedure:
  I.1.Take two spatula end full of lithium nitrate into a test tube and heat it strongly 
  until there is no further change. 
  2. Test the gases evolved with a damp blue litmus paper and a glowing splint.
  3.Observe and make conclusions on your observations.
  II .Repeat the procedure but using potassium nitrate/sodium nitrate
  

  Laboratory apparatus setting for thermal decomposition of a salt

  
  

  6.4.2. Heating the carbonates

  Activity 6.4.2

  Experiment: effect of heat on carbonates of group 1 elements
  In groups learners perform the following experiment, discuss and make conclusions 
  by explaining the observed phenomena, and write chemical reactions.
  Apparatus: glass test tubes, pair of tongs, Bunsen burner/heat source and spatula.
  Chemicals: lithium carbonate, calcium carbonate, potassium carbonate and 
  sodium carbonate, lime water
  Other requirements: match box.
  Procedure
  I.1 Take a spatula end full of lithium carbonate into a test tube and heat it strongly 
  until there is no further change. 
  

  2. Test the gases evolved with a damp blue litmus paper and lime water into another glass test tube as shown in the figure 

  

  3. Observe and make conclusions on your observations.
  II .Repeat the procedure but using calcium carbonate, potassium carbonate /
  sodium carbonate
  
  

  6.5. Solubility of group 1 compounds

  Activity 6.5
  a) Group 1 elements form ionic compounds.
  (i) Explain why.
  (ii)State the properties of ionic compounds
  b) Explain why lithium forms compounds with a covalent character contrarily to 
  other elements of the same group. State the properties of covalent compounds
  c) Both hydroxides and carbonates of lithium are less soluble than other hydroxides and carbonates of group 1.Why?
  

  

  

  Activity 6.6: 
  Experiment: Flame test of alkalis metals
  Materials: Mortar and pestle, beakers, Lithium carbonate, potassium sulphate, 
  sodium sulphate
  Procedure: Flame test wire /magnesia rod 
  NB: Wear your safety glasses.
  Dip the flame test wire/magnesia rod in the salt to be tested. Some of the salt 
  should stick to flame test wire/magnesia. 
  Gently wave the flame test wire/magnesia rod in the flame of the Bunsen 
  

  burner and note the color given of

  
  

  
  

  
  

  

  • NaOCl is used as bleaching agent and disinfectant
  • NaCl is used in seasoning food, preparing hydrogen chloride gas, in soap 
  production, manufacture of sodium, chlorine, sodium hydroxide and 
  sodium carbonate.
  • Molten sodium is used as a coolant in nuclear reactor. Its high thermal 
  conductivity and low melting temperature and the fact that its boiling 
  temperature is much higher than that of water make sodium suitable for 
  this purpose.
  • Sodium wire is used in electrical circuits for special applications. It is very 
  flexible and has a high electrical conductivity. The wire is coated with 
  plastics to exclude moisture.
  • Sodium vapor lamps are used for street lighting; the yellow light is 
  

  characteristic of sodium emission.

  

  

  

  

  

  

  a. The first member of the group often shows anomalous properties. Give two 
  properties in which the behavior of Lithium is abnormal and explain why.
  b. How does each of the following properties of the elements in Group 1 change 
  with the increasing atomic number? Explain why.
  i. Atomic radius
  ii. Ionization energy
  iii. Reducing properties
  iv. Reactivity with water
  v. Electronegativity
  c. How will successive ionization energies of Na vary?
  d. Why is the Na+ ion formed in normal chemical reaction rather than Na2+?
  

  e. How are ionization energies related to the reactivity of these elements?

  
  

• UNIT 7:TRENDS IN CHEMICAL PROPERTIES OF GROUP 2 ELEMENTS AND THEIR COMPOUNDS

  UNIT 7: TRENDS IN CHEMICAL PROPERTIES 
  OF GROUP 2 ELEMENTS AND THEIR 
  

  COMPOUNDS

  Key unit competence: compare and contrast the properties of the group 2 elements 

  and their compounds in relation to their position in the periodic table

  Learning objectives

  By the end of this unit, students should be able to:
  • Describe the physical properties of group 2 elements.
  • Describe the properties of group 2 oxides and hydroxides.
  • Explain the trends in the thermal decomposition of group 2 carbonates and 
  nitrates.
  • Explain the trends in the solubility of group 2 compounds.
  • State the uses of group 2 elements and their compounds.
  • Describe the industrial manufacture of cement. 
  • Discuss the environmental and health issues associated with the 
  manufacturing of the cement.
  • Perform experiments to compare and contrast the reactivity of group2 
  elements.
  • Write balanced equations of the reactions of group 2 elements, different 
  elements and the compounds.
  • Illustrate practically the trends in solubility and thermal decomposition of 
  group 2 compounds.
  • Test the alkaline character of group 2 hydroxides.
  • Be aware that the compounds of beryllium are different from the compounds 
  of the other group elements.
  • Perform chemical test for the presence of group 2 cations in solution.
  • Suggest preventive measures for environmental and health issues 
  associated with the manufacture of the cement.
  • Appreciate the logic underlying the position of elements in the periodic 
  table ,their electronic structure and the properties.
  • Appreciate the application of the chemistry of group 2 elements and their 
  compounds in the social economic development.
  • Develop the team work approach while performing experiment and writing 
  field –study reports. 
  
  • Develop the attitude of sustainable exploitation of natural resources .
  • Stimulate the culture of entrepreneurship in the area of chemistry.
  Introductory activity 7
  Complete the table below to identify the group 2 elements in the substances 
  

  found in our surroundings: at home and at the school.

  

  7.1. Occurrence and physical properties of group 2 elements
  Activities
  Name 1 or 2 elements of group 2 or their compounds that we commonly find in 
  

  Rwanda

  7.1.1. Occurrence
  Group 2 elements are active metals and are found in nature in form of compounds 
  or minerals such as: Limestone and marble for calcium, dolomite and magnesite for 
  magnesium etc… Hence Group 2 metals must be produced from the minerals they 
  

  are found in.

  

  7.1.2. Physical properties
  Group 2 elements are all metals, solid at room temperature; they are good conductors 
  of electricity. They have a silvery luster that soon disappears upon exposure to air. 
  They are malleable and ductile but less than alkali metals of Group 1. Their atomic 
  radius and their volume are smaller than those of Group 1 elements in the same 
  period. The Table 7.1 below shows some other physical properties of Group 2 
  elements.
  Atomic radius increases down the group due to increasing of electronic levels. 
  Melting and boiling temperature decreases down the group due to increasing of 
  atomic radius resulting in weakening of the metallic bond.
  The increasing of atomic radius explains also the decreasing of first ionization of 
  the elements down the group. This also explains that the metallic character of the 
  

  elements increases down the group.

  

  Checking up 7.1
  Question1: Metals are reducing agents because they lose easily electrons. You are given 
  3 elements of Group 2: Be, Ca and Ba. Which one are you going to choose as the best 
  reducing agent, and explain why?
  Question 2: How does each of the following properties of the elements in Group 2 
  change down the group and why?
  i) Atomic radius
  ii) Ionisation energy
  

  iii) Electropositivity

  

  7.2. Reactivity of group 2 elements
  Activities 7.2 (a)
  Activity 1:
  • Pour 200cm3
   of water in two different beakers
  • To the first beaker, add a small piece of magnesium ribbon. To the second 
  beaker, add a very small piece of sodium.
  • Record your observation.
  • Put a piece of blue and red litmus paper in both beakers
  • Record your observations.
  Activity 2:
  • Pour 200cm3 of water in pyrex beaker or borosilicate beaker 
  • Heat until water boils
  • Using crucible tongs, hold a large piece of magnesium in the steam
  

  • Record your observations

  

  

  

  

  7.3. Properties of group 2 compounds

  7.3.1. Ionic and covalent character of oxides and halides

  Activity 7.3.1 (a)
  • Pour 50 ml of paraffin in a beaker
  • Put 1g of calcium chloride and try to make a solution.
  • Pour 50 ml of water in another beaker
  • Put 1g of calcium chloride and try to make aqueous solution
  • Write down your observations and comments.
  Activity 7.3.1 (b)
  • Place a beaker on a table
  • Cut 15cm of magnesium ribbon
  • Using crucible tongs, hold and burn the magnesium ribbon over the 
  beaker.
  • What do you observe?
  • Add some water to the ash in the beaker.
  • Shake the mixture and add 2 drops of phenolphthalein or touch the 
  mixture with a red litmus paper
  

  • Record all your observations.

  
  

  
  

  
  

  
  

  
  

  
  

  The solubility of salts depends on two main opposite factors:
  The energy of dissociation of the crystal: The energy needed to dissociate the 
  solid crystal into its ions. This process requires energy; the process is endothermic.
  The energy of hydration of the ions produced: the amount of energy released 
  when ions undergo hydration or are surrounded by water molecules; this process is 
  exothermic.
  When the combination of the two processes above is in favor of the hydration of 
  ions, the salt is soluble; otherwise the salt is not soluble.
  The solubility will increase when the hydration process predominates more 
  

  and more the dissociation process and vice-versa.

  

  

  

  

  

  

  

  7.6. Uses of group 2 elements and their compounds
  Activity 7.6
  Describe the following compounds and show how each compound can be used 
  to prepare another if possible.
  a. Limestone b) Quicklime c)Slaked lime
  b. Have you heard about soil amendment in Rwanda? What is it?
  c. In groups, the students do research to find out how chalk used on 
  blackboard is produced.
  1. Beryllium
  Because beryllium is relatively light and has a wide temperature range, it can be 
  used in the manufacture of aircrafts’components.
  2. Magnesium
  • Chlorophyll, the pigment that absorbs light in plants, is a complex of 
  magnesiumand is necessary for photosynthesis. 
  • Magnesium hydroxide is used as Anti-acid medicine
  • Magnesium is used in making Grignard reagents, the organomagnesium 
  compounds.
  • Magnesium is used as sacrificial anode to prevent iron sheet from rusting.
  • Salts of magnesium and calcium are used in chemistry laboratory as drying 
  

  agents.

  

  

  

  

  

  

• UNIT8:TRENDS OF CHEMICAL PROPERTIES OF GROUP 13 ELEMENTS AND THEIR COMPOUNDS

  UNIT 8: TRENDS OF CHEMICAL PROPERTIES 
  OF GROUP 13 ELEMENTS AND THEIR 
  COMPOUNDS
  Key unit competence: Compare and contrast the chemical properties of the Group 13 
  elements and their compounds, in relation to their position in the Periodic Table.
  Learning objectives
  By the end of this unit, students should be able to:
  • State the physical properties of Group 13 elements
  • Explain the reactivity of Group 13 elements with oxygen, water, halogens, 
  dilute acids and sodium hydroxide
  • Describe the properties of oxides, hydroxides and chlorides of Group 13 
  elements
  • State the uses of Group 13 elements and their compounds
  • Compare and contrast the reactivity of Group 13 elements with oxygen, 
  water, halogens, dilute acids and sodium hydroxides
  • Perform experiments to show the solubility of Group 13 compounds
  • Practically illustrate the amphoteric properties of aluminium oxides and 
  hydoxides
  • Identify the anomalous properties of boron and its compounds
  • Perform chemical tests for the presence of aliminium ion in the solution.
  1. Consider the following elements: boron (z=5), aluminium (z=13), gallium (z=31), 
  indium (z=49), thallium (z=81) 
   (a)Write the electronic configuration of each element in term of s,p,d,f orbitals.
   (b) State the period and the block to which each element belongs.
  2. Draw diagram to show metallic bond in aluminium metal.
  3. Cite one known application of an element of Group 13 in our everyday life
  

  4. Explain why aluminium is a better conductor of electricity than sodium

  

  8.1. Physical properties of group 13 elements
  Activity 8.1
  In groups learners make research in library or on internet, discuss and explain why 
  a. Group 13 elements have higher melting point than group 1and 2 elements.
  b. Boron has higher ionization energy than other element of the same group.
  With the exception of boron, group 13 elements are metals. Boron is a non-metal 
  element with high melting point and low density.
  Aluminium is a metal element and has a low density, it is a good conductor of heat 
  and electricity, shiny, malleable, ductile and it has higher melting point than groups 
  1and 2 metals due to strong metallic bond resulting from 3 valency electrons 
  involved in making metallic bonding in aluminium metal.
  In small atoms electrons are held tightly and are difficult to remove, while in large 
  ones they are less tightly held since they are far away from the nucleus and are easy 
  to remove so that the ionization energy decreases down the group as the atomic 
  radius increases.
  The greater the forces of attraction and hence the boiling point and melting point 
  

  decrease down the group as the atomic radius increases.

  

  

  8.2. Reactions of aluminium

  8.2.1. Reaction of aluminium with oxygen

  
  

  
  

  
  

  
  

  8.2.3. Reaction with alkalis

  Activity 8.2 (d): 

  Reaction of aluminium with concentrated NaOH solution
  Experiment
  Learners perform experiments to investigate the reaction of aluminium with 
  NaOH solution
  Apparatuses: thermometer, pyrex beaker,stirrer
  Chemicals: aluminium powder,40% sodium hydroxide solution 
  Procedure:
  • Prepare 40% of sodium hydroxide by mixing 60cm3
   of water with 40g of 
  sodium hydroxide 
  • Take 0.5 g of aluminium powder into a pyrex beaker
  • Pour the solution of sodium hydroxide in the pyrex beaker containing 
  aluminium powder and allow the reaction to proceed for about 5 minutes.
  

  • Use thermometer to record the temperature during the process

  
  

  
  

  
  

  compounds; and only p electron will participate, hence the oxidation state +1.
  For lighter members such B and Al the s and p valency electrons, having almost the 
  same energy, are always available and used at the same time to form compounds 
  

  where they are in oxidation state +3.

  

  

  8.4. Anomalous properties of boron
  

  Activity 8.4
  

  In groups learners make research on internet and in library, discuss and explain 
  

  the following statement:
  

  a) Boron is a bad conductor of electricity
  

  b) Boron has higher boiling and melting points than other member of the group
  

  c) Boron oxide is an acidic oxide

  

  Checking up 8.4
  a) What is the cause of abnormal behavior of boron
  

  b) State any anomalous properties of boron

  
  

  
  

  8.6. Uses of some group 13 elements

  Activity 8.6

  a. Teacher brainstorms learners and ask them to talk about different applications 
  of aluminium and its compounds in daily life.
  b. Make research in libraries / internet and discuss about the use of aluminium,boron 
  and gallium and make presentations of your findings.
  Aluminium
  Aluminium is aboundant in the Earth’s crust and its applications are many. It is used 
  in:
  • making cooking utensils: this is because of its bright appearance and 
  lightness, resistance to corrosion, and its thermal conductivity
  • window frames or doors in buildings and houses
  • overhead high tension cables for distribution of electricity: this is because 
  of its low density and very good electrical conductivity.
  • alloys (e.g. Al and Mg) for the construction of airplanes and small boats due 
  to its lightness, malleability and higher tensile strength in the alloy.
  • Being completely resistant to corrosion it is ideal for packaging food
  • The insulating property of aluminium arises from its ability to reflect radiant 
  heat; this property is used in firefighters’ wear to reflect the heat from the 
  fire and keep them cool.
  • The polished surface of aluminium is used in the reflectors of car headlights
  • Aluminium is a component of clay (ibumba), mainly hydrous sulphate of 
  aluminium, used in the traditional production manufacture of clay pots 
  (ibibindi/ inkono).
  • Clay is also one of the basic raw materials in the production of Cement
  Boron
  Applications of boron are found in:
  • control rods to keep nuclear reactions in balance and avoid explosion; boron 
  absorbs excess neutrons preventing them from bombarding too many 
  uranium atoms which may result into explosion (fuel of nuclear reactors)
  • the manufacture of hard boron steel
  • as an additive to semiconductors silicon and germanium
  • the manufacture of borosilicate glass used in vacuum flasks and test tubes
  
  
  (a) Write a chemical equation to represent the reaction.
  (b) Why is it necessary to dry the chlorine?
  (c) What is the purpose of the soda lime?
  (d) Aluminium chloride is dissolved in water. Write the equation for the reaction that takes place
  2. (a) With reference to aluminium oxide, explain the term amphoteric oxide. Write 
  equations to illustrate.
  (b) Explain with chemical equations why aluminium utensils are not washed in 
  strong alkaline solutions.
  (c) Aluminium resists to corrosion. Comment and explain that popular saying.
  3. If you need to prepare aluminium hydroxide, why is it better to add a solution 
  of ammonia to a solution of aluminium salt, rather than to add a solution of 
  sodium hydroxide.
  4. How does gallium react with: 
  (a) hydrochloric acid
  (b) Sodium hydroxide
  5. Explain why aluminium is suitable for the following uses:
  (a) Manufacture of window frames
  (b)Electrical wiring
  (c)Packaging food
  (d)Suits for firefighters
  6. Water is suspected to contain calcium and aluminium ions. State a chemical test 
  that should be used to confirm the presence of the suspected ions. State the 
  reagent, observations and related chemical equation if any
  
  
  

  
  

  

• UNIT9:TRENDS OF CHEMICAL PROPERTIES OF GROUP 14 ELEMENTS AND THEIR COMPOUNDS

  Key unit Competence
  

  Compare and contrast the chemical properties of the Group 14 elements and their 
  
  

  compounds in relation to their position in the Periodic Table.
  

  Learning objectives
  

  By the end of this unit, students should be able to:
  

  • Compare and contrast the physical properties of Group 14 elements.
  

  • Compare the relative stabilities of the higher and lower oxidation states in 
  

  oxides.
  

  • Distinguish between the chemical reactions of the oxides and chlorides of 
  

  Group 14 elements.
  

  • Explain the trends in thermal stability of the oxide, halides and hydrides of 
  

  Group 14 elements.
  

  • Explain the variation in stability of oxidation state of +2 and +4 down the 
  

  Group 14 elements.
  

  Introductory Activity 9
  

  1. State any elements of group 14 that is found in Rwanda. Where are they produced from? What are they used for?
  

  2. State 2 allotropes of carbon and give a brief description of the structure of the 
  

  two allotropes.
  

  3. Explain the variation in electronegativity of group 13 elements as you move 
  

  down the group.
  

  4. Discuss the way the variation in size of atoms down a group affects their:
  

  a) Metallic character
  

  b) First ionization energy
  

   c) Ability to form ionic or covalent compounds.
  

  5. Describe the variation in melting points down group 1from lithium to potassium
  • Define the diagonal relationship.
  
  
  
  
  

  • Carbon, the first element of the group has two main allotropes: graphite 
  

  and diamond.
  

  • In graphite allotrope of carbon, each carbon is bonded to 3 other carbon 
  

  atoms to form a hexagonal structure. The structure of graphite is made of 
  

  hexagonal layers which are attracted to each other by weak Van der Waals 
  

  forces such that the layers slide over each other to make the structure 
  

  soft(Fig.9.1). In graphite structure, there are delocalised double bonds with 
  

  mobile electrons that allow graphite to conduct electricity.
  

  • In diamond, each carbon is covalently bonded to 4 other carbon atoms 
  

  forming a giant tetrahedral structure that makes it to be very hard (Fig.9.1). 
  

  In diamond, there are no mobile electrons as in graphite, hence diamond 
  

  does not conduct electricity.
  

  • As you move down the group in the carbon family, the atomic radius and 
  

  ionic radius increase while the electronegativity and ionization energy 
  

  decrease.
  

  • Atomic size increases on moving down the group due to additional 
  

  electronic shells.
  

  • Density increases as you move down the group.
  • Carbon is the only element in the family that can be found in pure form in 
  
  

  • Lead is the only element of group 14 that does not exist in various allotropes. 
  

  • Tin occurs as white, grey and rhombic tin.
  

  • Group 14 elements have much higher melting points and boiling points 
  

  than the group 13 elements.
  

  • Melting and boiling points tend to decrease as you move down the group 
  

  mainly because inter atomic bonding between the larger atoms reduce in 
  

  strength as you move down the group. 
  

  Moving down the group, there is an increase in atomic size which results in less 
  

  attraction of valence electrons by the nucleus. This change results in weaker metallic 
  

  bonding down the group and therefore there is a decrease in melting point, boiling 
  

  point, enthalpy change of atomization and first ionization energy.
  

  The decrease in first ionization energy from silicon to lead is relatively little compared 
  

  to that from carbon to silicon because there is a large increase in nuclear charge 
  

  which counterbalances the increase in atomic radius from silicon to lead. 
  ii) The increase in metallic character down the group causes a general increase in 
  

  conductivity. 
  

  Carbon is typically a solid, non-metal. Carbon graphite is a non-metal but conducts 
  

  electricity due to delocalized electrons in its structure. 
  

  In its compounds, carbon almost invariably completes its valence shell by forming 
  

  four covalent bonds
  

  Silicon is solid at room temperature and pressure, it is a semi-metallic element and 
  

  semi-conductor of electricity which is the second most abundant element on earth, 
  after oxygen. 
  
  
  It should also form bonds like C-C which are similar in strength to those of C and 
  other elements, particularly C-O bonds.
  Silicon forms -Si-O-Si-bonds predominantly. 
  ii) Multiple bonds
  Carbon forms double bonds and triple bonds between carbon atoms and that 
  bonding is formed by one Sigma bond and one π bond for double bond, one Sigma 
  

  bond and two π bonds in a triple bond.

  Checking-up 9.1
  1. Explain the reason why diamond has a higher melting point than silicon.
  2. Discuss the increase in metallic character when moving down in group 14 
  elements from carbon to lead.
  3. Diamond and graphite are allotropes of carbon,
  a) Draw their three dimensional structures.
  b) With reference to their structures, compare the hardness of diamond and 
  graphite.
  c) With reference to their structures, compare their electrical conductivity and 
  explain.
  4. Germanium has the same structure as diamond. Explain the type of bonds that 
  exist in the two elements.
  

  5. The first element in a group in the periodic table exhibits anomalous properties 
  compared with other members. Use carbon to illustrate this statement. 
  9.2. Chemical properties of Group 14 elements
  Activity 9.2 (a)
  1. Get a piece of charcoal and burn it. Observe and write the chemical equation 
  that represents the change that takes place when the charcoal burns.
  2. a) Put about 1 gram of carbon charcoal in a boiling tube.
  b) Add 1 ml of concentrated nitric acid.
  c) Heat strongly on a Bunsen burner flame using a test tube holder
  d) Observe and note the changes during heating.
  e) Deduce the chemical changes that have occurred.
  3. Write the molecular structure of carbon dioxide, carbonate ion and carbon 
  monoxide.
  4. Describe how CO2
   gas dissolves in water and state the nature of the solution 
  formed when it is in aqueous solution.
  5. Describe 2 chemical properties of amphoteric substances.
  
  
  
  
  Reaction of group 14 elements with acids and bases:
  

  Carbon does not react with dilute acids but reacts with hot, concentrated acids:

  

  

  

  9.3 Difference between the chemical reactions of the oxides 

  and chlorides of Group 14 elements.

  Activity 9.3
  1. Measure 0.5g of lead oxide or decompose the same quantity of lead nitrate 
  crystals by heating.
  2. Divide it into 5 portions and put each portion in a test tube.
  3. In the first test tube, add 2mL of dilute hydrochloric acid solution in which universal indicator has been dissolved.
  4. In the second test tube, add 2ml sodium hydroxide solution in which phenolphthalein indicator has been dissolved.
  

  5. Note the observations and deduce the acid–base nature of lead oxide.

  Interpretation of results of the above activity 
  

  The reactions that take place are:

  

  

  

  

   

  

  

  

  

  

  

  

  

  

  

  

  

  

  

  in chemical bonding in Ge, Sn and Pb elements of group 14 and hence only the 
  outermost p-electrons are involved.
  The electrons in s orbital are much more tightly bound to the nucleus than 
  p-electrons. As we move down the group, the difference in energy level between s 
  sub-shell and p sub-shell becomes wider. 
  So if we use weak oxidizing agents, only 2-p electrons are removed. If we use a strong 
  oxidizing agent 2 s-electrons and 2-p electrons are all removed from the shell. 
  If the elements in group 14 form +2 ions, they will lose the p electrons leaving the 
  s-electrons pair unused. For example, to form Pb2+ ions lead will lose the two 6p 
  electrons but the 6s electrons will remain in its sub-energy level.
  The inert pair effect shown in Pb2+ explains why the compounds of lead are 
  predominantly ionic
  
  
  Carbon uses:
  • As a component of fuel for combustion as charcoal or coal.
  • As the main component of crude oil and its derivatives used in our everyday 
  life such: fuel, plastics, etc…
  • As good chemical reducing agent used in extraction of metals (metallurgy).
  • As a lubricant in moving parts of machines, to make electrodes, in lead 
  pencils when mixed with clay. 
  • Carbon isotope, C-14 isotope is used in archaeological dating. 
  • Diamond is used to make glass cutters, drilling devices and as abrasive for 
  smoothing hard materials as precious gemstone in jewelry and ornamental 
  objects; it is also a precious stone appreciated in jewelry.
  Silicon uses:
  

  • Silicon is used as a semi-conductor in transistors in electrical gadgets such 
  as radios, computers, amplifiers etc..
  • Silicon in form of silicates is used in ceramics and in glass production.
  • Silicon is also used in medicine to make silicone implants.
  • Many rocks that we use for building our houses and other buildings are 
  Silicates.
  • Ferrosilicon alloy is used as a deoxidizer in steel manufacture.
  • Silicon dioxide can be used to produce toothpastes and in semiconductors; 
  silicon dioxide is the main component of sand, a raw material in the 
  manufacture of glass.
  Germanium uses:
  • Germanium being a metalloid, is used in transistors in electrical gadgets 
  such radios, computers, amplifiers etc..
  Tin uses:
  • Tin is used in plating steel sheets to resist corrosion; it is used for example 
  to make tinned cans to avoid the corrosion of the materials which are in 
  

  contact with an acid medium.

  

  

  9.7.1.The diagonal relationship in groups 1 & 2, 13 &14 elements
  Diagonal relationships are similarities between pairs of elements in different 
  groups which are adjacent to one another in the second and the third period of the 
  periodic table.
  These pairs are in Groups 1 and 2(Li/Mg), Groups 2 and 13(Be/Al) and Groups 13 
  and 14(B/Si). They exhibit similar properties; for example, boron and silicon are both 
  semi-conductors, they form halides that are hydrolyzed in water and have acidic 
  

  oxides.

  

  

  

  • Beryllium and aluminium have an appreciable covalent character of compounds 
  (e.g. the chlorides are predominantly covalent).
  9.7.4. Diagonal relationship between Boron and Silicon
  Due to its small size and similar charge/mass ratio, boron differs from other group 13 
  members, but it closely resembles silicon, the second element of group 14 to exhibit 
  diagonal relationship. Some important similarities between boron and silicon are 
  given below:
  • Both boron and silicon are typical non-metals that exist as non-metallic giant
  
  

  9.8. End unit assessment

  I: Fill in the following statements with a missing word:
  1. The arrangement of atoms in diamond structure is called………………..
  2…………………..is the only element of group 14 whose chloride does not 
  hydrolyse in water.
  

  3…………………is a semi-metallic element of group 14 whose oxide reacts with 
  HF acid only
  
  4…………………is the only element of group 14 that does not exist in various 
  allotropic forms.
  5………………….is the only element of group 14 whose compounds in the 
  oxidation state of +2 is more stable than that of +4.
  II. Answer the following questions:
  6. Write the equations for the reaction of decomposition of:
  a) Lead (II) hydroxide
  b) Tin tetrachloride
  7. Explain the amphoteric character of tinby using appropriate equations of 
  reaction.
  8. Discuss the stability of +2 oxidation state as you move down in group 14 
  elements.
  9. Explain the reason why the melting and boiling points of group 14 elements 
  

  decrease down the group.

  

  
  

• UNIT:TRENDS OF CHEMICAL PROPERTIES OF GROUP 15 ELEMENTS AND THEIR COMPOUNDS

  UNIT 10: TRENDS IN CHEMICAL PROPERTIES OF GROUP 
  15 ELEMENTS AND THEIR COMPOUNDS
  Key unit competency: Compare and contrast the properties of Group 15 elements 
  and their compounds, in relation to their position in the Periodic Table. 
  Learning objectives 
  By the end of this unit, students should be able to:
  • Describe the physical properties of Group 15 elements. 
  • Describe the variation in the metallic and non-metallic character of Group 
  15 elements. 
  • Explain recall the physical properties of the allotropes of phosphorus. 
  • Describe the chemical reactions of nitrogen compounds. 
  • Describe the impact of nitrogen oxides to the environment. 
  • Describe the industrial preparation of ammonia and nitric acid. 
  • Explain the reactions of nitric acid with metals and non-metals. 
  • Describe the chemical properties of phosphorus compounds. 
  

  • State the uses of the group 15 elements and its compounds

  

  

  10.1. Physical properties of group 15 elements and the 
  relative inertness of nitrogen
  Activity 10.1
  In pairs: 
  1. Assign the physical state for each of the elements in group 15 
  2. Explain what is meant by the term “metallic character”.
  3. Classify each element in this group as metal, non-metal or metalloid.
  4. Study the following figure carefully and answer the questions that follow
  
  a. Identify the molecule represented in the figure.
  b. What type of bond is there in the molecule?
  

  c. Suggest if the bond is strong or weak

  

  

  b. Metallic character
  Down group 15 elements, the atomic radius increases which makes the outermost 
  electron to be less attracted by the nucleus as you move down the group. Therefore, 
  less energy is required to remove the outermost electron, which results in the 
  increase in the metallic character down the group. This results also in decreasing of 
  ionization energy down the group.
  Nitrogen and phosphorous are non-metals, with the metallic properties first 
  appearing in arsenic and increasing down the group. Arsenic and antimony are 
  

  metalloids. Bismuth is a metal. 

  

  Checking Up 10.1
  1. Briefly describe how each of the following factors varies in group 15 elements:
  a) Atomic radius
  b) Electron affinity
  c) Melting point 
  d) First ionization energy
  2. Explain the following observations:
  a) In group 15 of the periodic table, metallic character increases as you move 
  down the group.
  b) The atomic radii of two elements A and B from group 15 are 0.121nm and 0.141nm 
  

  respectively. Identify the element with more metallic character. Justify your answer

  10.2. Reactions of group 15 elements

  

  All group 15 elements exhibit a common valency of three. They can complete their 
  octet structure in chemical combination by gaining three electrons.
  However, with the exception of nitrogen, group 15 elements have vacant d-orbitals 
  which they use to expand their octet to form compounds with a valency of five. For 
  instance phosphorous has a covalency of 5 due to availability of easily accessible 
  empty d orbitals which can be used for sp3
  d hybridization that allows it to have 5 
  

  unpaired electrons. Consider phosphorous, atomic number 15.

  

  

  

  

  

  c. i) State whether each oxide of A you have given in (b) is acidic, basic, neutral, or amphoteric and justify. 

  ii) Write the equation of reaction to illustrate your answer.

  10.3. Ammonia and nitric acid

  Activity 10.3 (a)

  Experiment: Laboratory preparation of ammonia 
  Materials and chemicals
  Round bottom flaskor hard glass test tube, U-tube, 3 corks, 10 grams of calcium 
  hydroxide, gas jars, bent delivery tube and straight delivery tube, 5 grams of 
  ammonium chloride on a watch glass and calcium oxide lumps.
  Procedure
  1. Set up the apparatus as shown in the diagram, with the chemicals indicated. Do 
  not start heating yet.
  2. When everything is in position, heat the hard flaskandcollect several gas jarsof 
  

  ammonia. Cover each jar with a glass slip and keep the jars for other experiments.

  
  

  Activity evaluation questions

  1. Record your observations
  2. Write a balanced equation of the reaction that take place.
  10.3.1. Laboratory preparation of ammonia and nitric acid
  a. Laboratory preparation of ammonia
  Ammonia is a covalent compound, consisting of nitrogen bonded to three hydrogen 
  atoms. It exists as a colourless gas at room temperature and it is naturally produced 
  during the decaying of nitrogenous organic compounds such as proteins. Ammonia 
  has a characteristic pungent odour.It is less dense than air and thus collected by 
  upward delivery method. In the laboratory it is prepared by heating a mixture of any 
  

  ammonium salt and an alkali.

  

  

  

  

  i. Uses of ammonia
  Agricultural industries are the major users of ammonia.  Ammonia and urea are used 
  as fertilizer, as very valuable source of nitrogen that is essential for plant growth.  
  Ammonia and urea are used as a source of protein in livestock feeds for ruminating 
  animals such as cattle, sheep and goats.  
  Ammonia can also be used as a pre-harvest cotton defoliant, an anti-fungal agent 
  on certain fruits and as preservative for the storage of high-moisture corn.
  The pulp and paper industry uses ammonia for pulping wood and as casein 
  dispersant in the coating of paper.
  The food and beverage industry uses ammonia as a source of nitrogen needed for 
  yeast and microorganisms involved in the fermentation process.
  ii. Environmental impact for industrial production of ammonia
  Making ammonia using the Haber process requires a lot of energy, which usually 
  involves burning fossil fuels. This releases carbon dioxide which causes global 
  warming. 
  b. Production of Nitric acid (Ostwald’s process)
  In the industrial manufacture of nitric acid a catalytic oxidation of ammonia to 
  nitrogen (II) oxide,NO, is carried out then a further oxidation of nitrogen (II) oxide 
  produces nitrogen (IV) oxide, NO2
  . Nitrogen dioxide is passed through water sprays 
  in a steel absorption tower to produce nitric acid. The excess nitrogen monoxide 
  is recycled back for more oxidation. Platinum is used as a catalyst. There are three 
  

  steps:

  

  

  

  

  

  

  

  

  

  

  

  b. Oxygen in the presence of a catalyst
  c. Copper (II) oxide
  d. Hydrochloric acid
  6. Write equations to show how nitric acid reacts with the following substances:
  a. Copper
  b. Sulphur
  

  c. Potassium hydroxide

  

  10.4.1. Allotropes of phosphorus
  By definition, allotropy is a property exhibited by some elements to exist in multiple 
  forms with different crystal structures. Allotropes are any two or more physical forms 
  in which an element can exist. Phosphorus exists in two main allotropic forms:
  

  • White phosphorus

  

  When prepared, ordinary phosphorus is white, but it turns light yellow when exposed 
  to sunlight. It is a crystalline, translucent, waxy solid, which glows faintly in moist 
  air and is extremely poisonous.It ignites spontaneously in air at 34°C and must be 
  stored under water. It is insoluble in water, slightly soluble in organic solvents, and 
  very soluble in carbon disulfide. White phosphorus melts at 44.1°C, boils at 280°C. 
  White phosphorus is prepared commercially by heating calcium phosphate with 
  sand (silicon dioxide) and coke in an electric furnace. When heated between 230°C 
  and 300°C in the absence of air, white phosphorus is converted into the red form.
  White phosphorus spontaneously takes fire in contact with air. White phosphorus is 
  

  considered and has been used as a chemical weapon.

  

  

  

  

  

  

  

  

  

  

  10.1.1 Environmental problems of using chemical fertilizers of nitrates 
  and phosphates
  Nitric acid is mainly used in the manufacture of nitrates fertilizers. Excess use of 
  nitrates as fertilizers is responsible of one type of pollutions of lakes and rivers called 
  eutrophication. 
  Eutrophication results from the excessive richness of nutrients in a lake or a water 
  body which causes a dense growth of plant life. When those water plants die and 
  are decomposed, during the decomposition process that uses oxygen, they deplete 
  the oxygen of the water body and render that water incapable of sustaining living 
  aquatic organisms. In that case, the body of water is said to be dead (biologically).
  The fraction of the nitrogen-based fertilizers which is not converted to be used 
  by plants accumulates in the soil or gets lost as run-off. High application rates of 
  nitrogen-containing fertilizers combined with the high water solubility of nitrate 
  leads to increased runoff into surface water as well as leaching into groundwater, 
  

  thereby causing groundwater pollution.

  

  

  

  

  
  

• UNIT11:TRENDS OF CHEMICAL PROPERTIES OF GROUP 16 ELEMENTS AND THEIR COMPOUNDS

  UNIT 11: TRENDS OF CHEMICAL PROPERTIES OF GROUP 
  16 ELEMENTS AND THEIR COMPOUNDS
  Key unit competence: 
  Compare and contrast the chemical properties of the Group 16 elements and their 
  compounds in relation to their position in the Periodic Table.
  Learning objectives
  By the end of this unit, students should be able to:
  • Describe the physical properties of Group 16 elements. 
  • Describe the reactions between sulphur and oxygen.
  • Describe the steps and conditions applied in the industrial preparations of 
  sulphuric acid. 
  • Describe the chemical properties of sulphuric acid. 
  • Describe the properties of oxoanions. 
  

  • State uses of the Group 16 elements and compounds.

  

  

  

  

  

  

  

  

  

  11.2. Comparison of acidity and volatility of group 16 
  hydrides
  Activity 11.2 
  1. In pairs, carry out research and write a note on the following terms:
  a. Hydrides 
  b. The strength of an acid
  d. A weak acid
  e. A strong acid
  2. With an example, explain what is meant by the term “hydrogen bond” and show 
  

  how it is formed.

  

  

  

  Safety: 
  

  Sulfuric acid is a very strong acid and is extremely corrosive to skin. Wear gloves 
  

  and safety goggles. During the reaction, steam is generated. It is hot. It is 
  recommended to work in a fume cupboard.
  Procedure:
  Spread some paper towels on the tray.
  1. Put sugar into 300 ml beaker. 
  2. Insert stirring rod into center of sugar.
  3. Put beaker on paper towels on the tray.
  4. Add 70 ml of sulfuric acid to the sugar and stir briefly.
  5. Stand about 1 - 2 meters away and wait for reaction to begin and observe what 
  will happen.
  Clean Up: You might want to incorporate part of the clean up procedure into the 
  demonstration. 
  Remove black carbon column from the beaker and put it into a liter beaker with 
  some sodium bicarbonate (hydrogen carbonate). With spatula, break the column 
  of carbon into smaller pieces. Add a little water and set back on the tray. The 
  foaming action is also exciting.
  Neutralize any acid spills with sodium bicarbonate and wipe clean. Leave lecture 
  hall clean for the next class.
  Rinse all glassware and carbon chunks with lots of water. Carbon can be thrown 
  away in trash.
  Study questions
  1. Record your observations
  

  2. Write an equation for a reaction that takes place in this experiment

  

  

  Step 1: Production of sulphur dioxide
  Sulphur dioxide is obtained by either burning elementary sulphur or roasting metal 
  

  sulphides in air in combustion chamber.

  

  

  

  Checking Up 11.3
  1. a) Describe the Haber or Contact process for the manufacture of sulphuric acid.
  b) Why is sulphur trioxide formed in this process not absorbed directly in water?
  2. Concentrated sulphuric acid acts as a dehydrating agent. What does it mean? 
  3. Write equations to show how concentrated sulphuric acid reacts with:
  a. Zinc
  b. Magnesium
  

  c. Carbon

  11.4 Properties of oxoanions of sulphur
  Activity 11.4 (a)
  1. Use the library and/or internet to explain the following:
  a. Oxidation 
  i) In terms of oxidation state 
  ii) In terms of electron transfer
  b. Reduction
  i) In terms of oxidation state
  ii) In terms of electron transfer
  c. Oxidizing agent
  d. Reducing agent
  Activity 11.4 (b)
  2. An experiment for Heating hydrated copper(II) sulfate
  Objectives: 
  Students remove the water of crystallisation from hydrated copper (II) sulfate 
  by heating. Condensing in a test-tube collects the water. The white anhydrous 
  copper (II) sulfate can then be rehydrated, the blue colour returns.
  Apparatus and equipment (per group) 
  

  1. Two test-tubes

  

  

  

  potassium sulphate, and calcium sulphate are not decomposed by heat.
  Only certain sulphate salts are decomposed by heat when heated strongly. On 
  heating, some sulphates decompose to give either sulphur trioxide or sulphur 
  

  dioxide or both.

  

  Checking Up 11.4
  1. Write equations to show how thiosulfate ions reduce the following substances:
  a. Iodine
  b. Iron (III) ion
  c. Aluminium ion
  2. Write equations to show the action of heat on the following sulphates:
  a. Zinc (II) sulphate
  b. Iron (III) sulphate
  c. Copper (II) sulphate
  3. When hydrated copper II sulphate solid is heated in a boiling tube, a white solid 
  Q and droplets of a colourless liquid P are observed.
  a. Identify substances; liquid P and solid Q.
  b. Explain the observation above.
  Explain what would be observed if water is added to white solid Q.
  11.5 Identification of sulphite and sulphate ions
  Activity 11.5
  Given a substance Y which contains one cation and one anion, identify the cation 
  and the anion in Y. Carry out the following tests on Y and record your observations 
  

  and deductions in the table below.

  

  

  

   

  

  

  11.6.1.Uses of oxygen
  The first use of oxygen is in breathing and metabolism processes of all living 
  organisms.
  There are many other commercial uses for oxygen gas, which is typically obtained 
  through fractional distillation of air. It is used in all operations involving combustion 
  as the active component of air.
  It is used in the manufacture of iron, steel, and other chemicals. Oxygen is also used 
  as an oxidizer in rocket fuel, and for medicinal purposes. Mixture of oxygen and 
  ethyne (oxyacetylene) is used for welding and metal cutting.
  11.6.2. Uses of sulphur
  The main use of Sulphur is the manufacture of sulphuric acid.
  Sulphur is also used in vulcanization of rubber, a  chemical process  for 
  converting  natural rubber  or related  polymers  into more durable and pressure 
  resisting materials by heating them with  sulfur  or other equivalent curatives 
  or accelerators. These additives modify the polymer by forming cross-links (bridges) 
  between individual polymer chains, making the final product very hard and resistant 
  to pressure and other conditions.
  Sulphur is an ingredient in the manufacture of dyes, fireworks and other sulphur 
  compounds.
  11.6.3. Uses of sulphuric acid
  Sulphuric acid is a very important industrial chemical. It used to be called the giant 
  of chemical industry. It is used in the manufacture of hundreds of other compounds 
  in many industrial processes. 
  • The bulk of sulphuric acid produced is used in the manufacture of fertilisers 
  (e.g., ammonium sulphate, superphosphate). 
  • Sulfuric acid is also used in many other applications such as in: metallurgical 
  

  industry, storage batteries, chemistry laboratories, etc….

  

  

  

• UNIT12:TRENDS OF CHEMICAL PROPERTIES OF GROUP 17 ELEMENTS AND THEIR COMPOUNDS

  UNIT 12: TRENDS OF CHEMICAL PROPERTIES OF GROUP 
  17 ELEMENTS AND THEIR COMPOUNDS
  Key unit Competence:
  Compare and contrast the chemical properties of the Group 17 elements and their 
  compounds in relation to their position in the Periodic Table.
  Learning objectives
  By the end of this unit, students should be able to:
  • Prepare and test halogens.
  • Perform experiments to prepare and test chlorine, bromine and iodine.
  • Relate the oxidizing power of Group17 elements to their reactivity.
  • Relate the acidity strength of oxoacids to the number of oxygen atoms 
  combined with the halogen.
  • Compare the reactions of the halogens with cold dilute sodium hydroxide 
  and hot concentrated sodium hydroxide solutions.
  • State the uses and hazards of halogens and their compounds.
  • Test for the presence of halides ions in aqueous solutions
  • State the natural occurrence of halogens
  • Describe the extraction methods of halogens
  • Explain the trends of physical and chemical properties of Group 17 elements 
  down the group
  • Describe the trends in strength acidity, volatility and reducing power of 
  halogens hydrides
  • Describe the chemical properties of chlorates, iodates, perchlorates and 
  

  periodates

  

  

  Checking up 12.1
  1. State 2 locations where chlorine can be found in nature.
  2. Write the chemical formulae of the compounds of halogens in nature.
  3. Give the name of one lake in Rwanda where salt is abundant in water. 
  4. Explain the separation method you can use to get the salt crystals from the 
  water. 
  5. Fluoride ion is the most difficult to oxidise into fluorine, whereas iodide ion is 
  

  the easiest. Explain why.

  12.2. Preparation methods of halogens

  
  

  
  

  Repeat procedures steps 1 to 7 but this time use solutions of KBr or NaBr instead 
  of KI in the boiling tube to prepare bromine and chlorine.
  Activity 12.2 (b) Preparation of bromine and iodine
  1. Put 0.5 gram of MnO2
   in a round bottomed flask.
  2. Pour concentrated NaBr solution (5 ml of a 0.1 mol/litre) in the round bottomed 
  flask.
  3. Pour 5 ml of 1 mol/litre HCl solution in the round bottomed flask mixture.
  4. Connect the apparatus to a delivery tube using a rubber stopper.
  5. Heat the round bottomed flask mixture.
  6. Direct the delivery tube in a solution of KI in a test tube.
  

  7. Note the observable changes.

  

  Activity 12.2. (c): Electrolysis of concentrated NaCl solution
  a) Put 1 g of NaCl in a beaker.
  b) Add water and stir using a glass rod until all the salt dissolves.
  c) Pour the solution in an electrolyser.
  

  d) Connect the electrolyser to the source of direct current and switch on.
  

  e) Dip a test tube full of water in the NaCl solution in inverted position from above 
  each electrode
  f) Put 2 drops of phenolphthalein indicator in the solution under each test tube.
  g) Record the observations that take place for 5 minutes.
  

  Apparatus set-up: Electrolysis of concentrated NaCl solution

  

  12.2.1. Chlorine
  Most commercial chlorine is obtained by electrolysis of chloride ions in aqueous 
  solutions of sodium chloride or molten NaCl.
  

  The reactions that take place are shown by the following chemical equations: 

  

  

  

  

  

  Interpretations
  Chlorine is liberated by the reaction between 2M hydrochloric acid and potassium 
  permanganate solution. Chlorine displaces iodine from potassium iodide solution, 
  

  which dissolves in water to give a dark-red solution, and turns starch indicator dark blue. The greyish-black residue is due to the formation of Iodine solid.

  

  

   

  

  

  

  

  

  

  

  12.4. Preparation of Hydrogen halides
  Activity 12.4.(a)
  Laboratory preparation of Chlorine
  Reactants:Sodium Chloride and sulfuric acid
  Rocedure:
  1.Put 50g of NaCl in round bottomed flask
  2.Pour conc sulfuric acid throuth the filter fannel and heat
  3.The liberated gas is pased throuth the concentrated sulfuric acid
  4.Collect the gas by the Downward delivery
  Note:The lower end of the thistle funnel must be dippen in acid,or you can use the funnel 
  

  with Syphon

  

  

  

  

  12.5. Trends in strength of acidity, volatility and reducing 
  power of hydrogen halides
  12.5.1. Acid strength
  The acid strength is a measure of how an acid dissociates in water into its ions. 
  Strong acids dissociate completely into their ions, whereas a weak acid dissociates 
  

  partially into its ions.

  

  

  bonds. It is liquid at room temperature while other hydrogen halides are gases. 
  

  The trend in volatility:

  

  

  

  12.5.4. Tests for halide ions in aqueous solution
  

  Test of substance X with an unknown anion

  Activity 12.5 

  Identification of ions:

  i) You are provided with a solution of X substance. 
  ii) Put 1 ml of X solution in each of the 4 test tubes.
  iii) Add in each test tube the reagent solutions as indicated in the table below. 
  

  iv) Note down the observations for interpretation later in each test.

  
  

  
  

  
  

  
  

  12.6. Chemical properties of chlorates, iodates, perchlorates 
  

  and periodates

  
  

  
  

  
  

  
  

  
  

  
  

  12.7. Uses of halogens and their compounds

  Activity 12.7

  1. When you want to eat food, salt is dissolved in it. Indicate the chemical composition 
  of table salt and its natural occurrence. 
  2. Chlorine compounds are used in the treatment of water. Explain how chlorine reacts to be a good disinfectant in water treatment.
   3. a)Write the observations of the phenomenon that takes place when electrolysis of 
  a concentrated solution of chlorine is carried out in the laboratory. b) Deduce the 
  product of reaction that is formed at the anode.
  12.7.1. Uses of Halogens
  

  Halogens and their compounds have many applications and uses in different 

  
  

  12.7.2. Hazards caused by group 17 elements
  Bromine effects
  • On heating, toxic fumes are formed. 
  • Reacts violently (explosively) with many compounds. 
  • Attacks plastics, rubber and coatings. 
  Chlorine effects
  • It reacts violently with many compounds like ammonia and may cause fire 
  and explosion. 
  • It attacks many metals in the presence of water. 
  • It attacks plastics, rubber and coatings.
  Chlorine oxide effects
  • It may explosively decompose when it encounters shock and friction then 
  it may explode on heating. 
  • It reacts violently with mercury, phosphorus, sulphur, etc causing fire and 
  explosion hazard.
  

  Fluorine effects
  

  • It reacts violently with water to produce toxic and corrosive vapours: ozone 
  and hydrogen fluoride. 
  • It reacts violently with ammonia, metals, oxidants, etc, to cause fire and 
  explosion.
  Hydrogen bromide effects
  • It reacts violently with strong oxidants and many organic compounds to 
  cause fire and explosion. 
  • It attacks many metals forming flammable hydrogen gas.
  Hydrogen fluoride effects
  • It reacts violently with many compounds causing fire and explosion. 
  • On contact with air, it emits corrosive fumes which are heavier than air. 
  • It attacks glass and other silicon-containing compounds.
  SF6 effects
  • The substance decomposes in a fire to produce toxic fumes of sulphur 
  oxides and hydrogen fluoride 
  • When it is heated, there is formation of toxic fumes.
  
  12.8. End of unit assessment
  
  
  
  
  

  

  

  

• UNIT 13:PROPERTIES AND USES OF GROUP 18 ELEMENTS AND THEIR COMPOUNDS

  UNIT 13: PROPERTIES AND USES OF GROUP 18 
  ELEMENTS AND THEIR COMPOUNDS
  Key unit competence: Compare and contrast the properties of the group 18 
  elements in relation to their position in the periodic table.
  Learning objectives
  By the end of this unit, students should be able to:
  • State the physical properties of the Group 18 elements.
  • Explain the lack of reactivity of the group 18 elements.
  • Associate chemical inertia of the group 18 elements to their full valence 
  shell.
  

  • Recognize the importance of noble gases or group 18 elements in the daily 
  life.
  Introductory activity
  Make a research to find out the type gas :
  Inside the Bulb, in balloon, responsible for different colors dispayed by this house 
  

  (or in advertising sings)

  

  

  13.1. Occurrence and physical properties of noble gases

  Activity 13.1

  The air is composed of a mixture of gases including water vapour.

  i)Make a research (with any documentation) to identify its components and arrange them according to their abundances (Component1> Component 2, 
  etc…)
  ii)Show how these components can react each other if possible 
  • If not possible, justify your answer.
  iii)Explain how neon lamp works
  13.1.1. Occurrence
  • All the noble gases except radon occur in the atmosphere. Their total 
  atmospheric abundance in air is 0.03%; argon is the major component.
  • Helium and sometimes neon are found in minerals of radioactive origin 
  

  e.g., pitchblende, monazite, cleveite.

  

  

  

  Checking up 13.2
  Question: Explain why in some applications such as air balloons, helium is 
  preferred to hydrogen? 
  13.3. Uses of noble gases
  Activity: 13.3 
  Do a research (with any documentation) to find how each noble gas has been 
  discovered and its uses? 
  

  Helium
  • Helium is a non-inflammable and light gas. Hence, it is used in filling balloons 
  for meteorological observations, replacing the flammable hydrogen gas.
  • It is also used in gas-cooled nuclear reactors. 
  • Liquid helium (B.P:-267.8o
  C) finds use as cryogenic agent for carrying out 
  various experiments and conservation at very low temperatures.
  Neon
  • Neon is used in advertising signs, it glows when electricity is passed through 
  it. Different coloured neon lights can be made by coating the inside of the 
  glass tubes with colored chemicals.
  

  • Neon bulbs are more used in our daily life. 

  

  Argon
  • It is used in light bulbs. The very thin metal filament inside the bulb would 
  react with oxygen and burn away if the bulb were filled with air instead of 
  argon. 
  • Argon is used mainly to provide an inert atmosphere in high temperature 
  metallurgical processes (arc welding of metals or alloys).
  • It is also used in the laboratory for handling substances that are air-sensitive.
  Krypton
  Krypton is used in lasers. Krypton lasers are used by surgeons to treat certain eye 
  problems. It is used in light bulbs designed for special purposes.
  Xenon
  Xenon is used in fluorescent bulbs, flash bulbs and lasers. Xenon emits an instant, 
  intense light when present in discharge tubes. This property of xenon is utilized in 
  high-speed electronic flash bulbs used by photographers.
  Radon
  Radon is radioactive and is used in medicine as a source of gamma rays. The gas is 
  sealed in small capsules, which are implanted in the body to destroy malignant (e.g., 
  

  cancerous) growths.

  

  13.4. End unit assessment
  1. a) Give a reason why the first ionization energies of noble gases are very high.
  b) State one use of neon and give a reason to support your answer.
  c) State and explain the trend in atomic radius among noble gases.
  d) Why are noble gases unreactive?
  e) Explain why the value of the first ionisation energy of neon is higher than that 
  of sodium.
  2. Explain why Group 18 elements are rare on Earth?
  3. The discovery of compounds of noble gases has been done, up to date, with Xe 
  

  and Kr, not with He or Ne. Can you suggest a probable reason?

• UNIT14:TRENDS OF CHEMICAL PROPERTIES OF PERIOD 3 ELEMENTS AND THEIR COMPOUNDS

  UNIT 14: TRENDS IN CHEMICAL PROPERTIES OF PERIOD 3 
  ELEMENTS AND THEIR COMPOUNDS
  Key unit competency: Compare and contrast the properties of the Period 3 elements and 
  their compounds in relation to their positions in the Periodic Table. 
  Learning objectives:
  By the end of this unit, students should be able to:
  • Compare the physical properties of the Period 3 elements. 
  • Describe the nature of the oxides of the Period 3 elements and the type of 
  bonding in their chlorides, oxides and hydrides. 
  • Relate the physical properties of the Period 3 elements to their position in 
  Periodic Table. 
  • Relate the physical properties of compounds of the Period 3 elements to 
  

  their nature of bonds across the period.

  

  

  14.1. Physical Properties of the Period 3 elements
  Activity 14.1
  1. Write the electronic configuration of the following elements in terms of s, p, d 
  and f…
  (i) Sodium (ii) Magnesium (iii) Aluminium (iv) phosphorous (v) sulphur
  2. Considering the electronic configuration of magnesium and Aluminium, phosphorus and sulphur. How do you expect their ionization energies to vary?
  3. How do you expect the general trend in ionization energy, electron affinity, 
  melting and boiling point, electronegativity to vary for the elements in the period 3?
  4. Considering the electronic configuration of magnesium and Aluminium, phosphorus and sulphur. What can you say about them, how do you expect their 
  ionization energies to vary?
  (a) Variation of First ionization energies (IE) of Period 3 elements
  First ionization energy generally increases across Period 3 from left to right. However, 
  it drops at aluminium and Sulphur (table 14.1 and Fig.14.1). This can be explained in 
  

  term of more stable electronic structures of the two elements after losing 1 electron:

  

  

  

  Going across Period 3 from left to right, the number of protons in the nucleus 
  increases so, the nuclear charge increases. There are more electrons, but the 
  increase in shielding is negligible because each extra electron enters the same 
  principal energy level. Therefore, the force of attraction between the nucleus and 
  the electrons increases. So the atomic radius decreases as indicated in the Figure 
  14.2 and table 14.2. 
  (c)Variation of electronegativity of Period 3 elements
  

  Table 14.3: Variation of electronegativity of period 3 elements

  
  

  Figure 14.3: Graph showing the variation of electronegativity of period 3 elements
  
  Going across Period 3 from left to right, electronegativity increases almost linearly 
  due to the nuclear charge increase as atomic radius decreases. There are more 
  electrons, but the increase in shielding is negligible because each extra electron 
  enters the same principal energy level so electrons will be more strongly attracted 
  to the nucleus.
  You might expect argon (with 18 electrons) to be the most electronegative element 
  in Period 3, but its outermost energy level is full. Therefore, it does not form covalent 
  bonds with other atoms, so it is given an electronegativity value of zero.
  d. Variation of melting and boiling points in Period 3
  Melting and boiling points generally increase going from sodium to silicon, then 
  

  decrease going to argon with a “jump” at Sulphur (Fig 14.4 and Table 14.4).

  
  

  
  

  
  

  
  

  The delocalized electrons are free to move and carry charge. Going from sodium to 
  aluminium, the number of delocalized electrons increases, there are more electrons 
  which can move and carry charge so the electrical conductivity increases.
  Silicon is called a semi-conductor because at higher temperatures more electrons 
  are promoted to the higher energy levels so there are more delocalized electrons to 
  move and carry charge.
  Phosphorus, sulphur and chlorine, the outer electrons are not free to move and carry 
  charge because they are held strongly in covalent bonds. In argon (mono atomic) the 
  outer electrons are not free to move and carry charge because they are held strongly 
  in a stable third energy level and this explains their zero electrical conductivity.
  f. Variation of metallic character of period 3 elements
  Metallic character decreases as you move across a period 3 in the periodic table from 
  left to right. This occurs as atoms more readily accept electrons to fill the valence 
  shell than lose them. Note that as the metallic character decreases across the period, 
  the reducing power decreases whereas oxidizing power increases. 
  g. Variation of electron affinity across period 3 elements
  The electron affinity [EA] is the energy change for the process of adding an electron 
  

  to a gaseous atom to form an anion (negative ion).

  

  Checking up 14.1
  1.Explain the variation of the following terms as applied in period 3 of the periodic 
  table:
  (i) Ionization energy, Electronegativity, )
  (ii)Explain the anomalous behavior indicated by magnesium and phosphorous in graph 14.1 above 
  2.The table below shows the melting points of the period 3 elements except for 
  

  silicon:

  

  (a)Explain in terms of bonding why the melting point of magnesium is higher 
  than that of sodium.
   (b) Predict the approximate melting point of silicon.
  (c) Explain why chlorine has a lower melting point than sulphur.
  (d) Explain the variation of metallic character, electronegativity, atomic radii 
  ,first ionization energy, melting and boiling points, electron affinity and 
  

  electrical conductivity across the period

  14.2. Chemical properties of period 3 elements

  
  

  Study questions:

  1. What do you say about your observations made in experiment above.
  2. Write equation for the reaction that occurs in each test tube in procedure 2.
  (b) Experiment to investigate the action of heat on period 3 elements
  Materials /apparatus:
  Water , test tubes, a piece of sodium metal, aluminium power/sheet, magnesium 
  ribbon/powder, phosphorous and sulphur powder, universal indicator , pair of 
  tongs, source of heat
  Procedure: 
  1. Hold a piece of magnesium ribbon on a Bunsen flame and record you observation.
  2. Repeat experiment 1 for sodium, aluminium, phosphorous and sulphur and 
  record your observation in each case.
  3. For each of the products formed i.e. for metal oxides formed, add water and dip 
  a litmus paper to test their nature. 
  Note: if the oxide is gaseous hold a piece of litmus paper on the mouth of the test tube.
  Study questions:
  1.Write equations to show how the metals react with oxygen.
  2. What would you expect to observe when the metal is burned in oxygen.
  a) Reaction with water 
  Reactivity with water generally decreases across the period from left to right because 
  there is a decrease in metallic properties.
  i) Sodium reacts vigorously with cold water to form sodium hydroxide and hydrogen 
  

  gas. 

  

  

  

  

  

  

  14.3. Compound of period 3 elements
  The oxides of period 3 elements:
  Activity 14.3(a)
  1. Write the formulae of the oxides of period 3 elements
  2. What did you consider when writing the formulae of the oxide in 1 above?
  3. How do you expect the oxides to behave in water? Explain your answer.
  

  4. Suggest the trend of acid- base character of the oxides of period 3

  

  middle, silicon forms a giant covalent oxide (silicon dioxide); the elements on the 
  right form simple molecular oxides with simple structures. The intermolecular forces 
  binding one molecule to its neighbors are van der Waals dispersion forces or dipoledipole interactions. 
  Physical properties of the oxides of period 3 elements
  Melting and boiling points: the metal oxides and silicon dioxide have high melting and 
  boiling points because a large amount of energy is needed to break the strong bonds (ionic 
  or covalent) operating in three dimensions. The oxides of phosphorus, sulfur and chlorine 
  consist of individual molecules.
  Electrical conductivity: None of the  oxides above have any free or mobile electrons, 
  

  indicating that none of them will conduct electricity when solid. 

  

  

  Acid-base Behavior of the Oxides
  Activity 14.3 (b)
  1. Classify the oxides in terms physical states of the oxides of period 3.
  2. How do you expect the oxides react with water, acids, and sodium hydroxide.
  (use equations to justify your answer)
  3.(a) Predict the nature of oxides of period 3 elements when dissolved in water.
   (b)What would you expect to observe when both blue and red litmus papers 
  are dropped into each of the solutions formed in question (2) above in water.
  Acidity increases from left to right, ranging from strongly basic oxides on the left to 
  acidic ones on the right, with an amphoteric oxide (aluminum oxide) in the middle. 
  

  Reaction of oxides with water: 

  

  

  

  

  
  

  
  

  
  

  
  

  
  

  Electrical conductivity: solid chlorides do not conduct electricity because the ions 
  are not free to move
  Sodium, magnesium and aluminium chlorides are ionic and so will conduct 
  electricity when they are molten or in aqueous solution. The rest of the chlorides do 
  

  not conduct either in solution or molten state due to absence of ions

  Reactions with water

  
  

  Checking up 14.3(c)

  1. a. Distinguish between dissolving and hydrolysis.
  

  b. Name one chloride that dissolves in water, and one chloride that undergo 
  

  hydrolysis.
  

  c. State how the bonding in the chlorides changes on crossing the second and 
  third periods from left to right
  Checking up 14.3(c)
  1. a. Distinguish between dissolving and hydrolysis.
  b. Name one chloride that dissolves in water, and one chloride that undergo 
  hydrolysis.
  c. State how the bonding in the chlorides changes on crossing the second and 
  

  third periods from left to right

  14.3.3. The hydrides of period 3 elements
  Activity 14.4(d)
  1. Period 3 elements from sodium to chlorine form different hydrides of different 
  bond nature, physical properties and structure.
  (a) Write the formula of the hydrides formed by period 3 elements.
  (b) Predict the nature of bonding based on your knowledge of periodicity of elements in the periodic table.
  (c ) Basing on the nature of bonding predicted in (b) above. How would you expect their boiling and melting point vary across the period?
  (d) Predict the nature of solutions formed by hydrides when dissolved in water. 
  What would you expect to observe if red and blue litmus papers were separately dropped into each solution?
  Hydrides are commonly named after binary compounds that hydrogen forms with 
  other elements of the periodic table. Hydride compounds in general form with 
  almost any element, except a few noble gases. The common hydrides of period 3 
  

  elements are as shown in the table 14.6 below

  

  

  The hydrides above are examples of period 3 elements with some of their properties 
  summarized in the tabl. As we can see the hydrides of period 3 vary from ionic 
  hydride such as NaH at the left side to polar covalent hydride such as HCl at the right 
  

  side of the period.

  

  

  

  

  a. In terms of crystal structure and bonding, explain in each case why the melting 
  points of sodium oxide and silicon dioxide are high.
  b. Predict whether the melting point of lithium oxide is higher than, the same as, 
  or lower than the melting point of sodium oxide and explain your prediction.
  c. Phosphorus (V) oxide has a lower melting point than sodium oxide.
  i. State the structure of and bonding in phosphorus (V) oxide.
  ii.  Explain why the melting point of phosphorus(V) oxide is low.
  d. Samples of phosphorus(V) oxide and sodium oxide were reacted with water.
  In each case, predict the pH of the solution formed and write an equation for 
  the reaction.
  4. Sodium chloride is a high melting point solid which dissolves in water to make 
  a colorless solution. Silicon (IV) chloride is a liquid at room temperature which 
  fumes in moist air, and reacts violently with water.
  a. Draw a diagram to show the arrangement of the particles in solid sodium 
  chloride, making clear exactly what particles you are talking about.
  b. Explain why this arrangement leads to a high melting point.
  c. Draw a simple diagram to show the structure of silicon (IV) chloride, and explain 
  why silicon (IV) chloride is a liquid at room temperature.
  d. Why is there such a big difference between the chlorides of sodium and silicon?
  e. Briefly describe and explain the difference in electrical conductivity between 
  sodium chloride and silicon (IV) chloride in both solid and aqueous molten 
  state.
  f. Write an equation to show what happens when silicon (IV) chloride reacts with 
  water.
  g. Name another Period 3 chloride which behaves similarly to sodium chloride, 
  and one which behaves similarly to silicon (IV) chloride.
  5. With the help of equation describe how the hydrides of period 3 react with 
  

  water.

  14.4. End unit assessment
  

  1. Use the information in the following table to explain the statements below

  

  

  3. The elements Sodium, Magnesium, silicon, phosphorous and chlorine are 
  members of the third period of the periodic table
  a. i. Write down the formula of the principal oxides and chlorides of the elements 
  listed above and in each case indicate the type of bonding.
  ii. Explain what happens when each of the above oxides and chloride is added to 
  water and indicate whether the resultant solution will be acidic, basic or neutral.
  c. The melting points of Mg , Si and S are 6500
  C, 14230
  C respectively. Explain the 
  differences in the melting points of the elements.
  d. Name the type of bonding that exists in the hydrides of the elements Sodium, 
  Phosphorous and sulphur and write the equations to show the reactions if any 
  of the hydrides with water.
  4. Choose from the elements: Sodium, magnesium aluminium, silicon, phosphorous, chorine and argon 
  a. List the elements that react readily with cold water to form alkaline solutions. 
  And write the equations for the reactions.
  b. List the hydrides that have hydrides with low boiling points/temperatures and 
  explain why.
  c. List the elements that form nitrates and write the formulae of nitrates.
  d. What is the most ionic compound that can be formed by the combination of 
  two of these elements.
  e. Which element has both metallic and non metallic properties?
  f. Name the elements that normally exist as molecules
  
  
  

• UNIT15:FACTORS THAT AFFECT CHEMICAL EQUILIBRIUM

  UNIT15: FACTORS THAT AFFECT CHEMICAL EQUILIBRIUM 
  Key unit competency: Deduce how concentration, pressure, catalyst and 
  

  temperature affect the chemical processes in industry.
  

  
  

  Learning Objectives
  By the end of this unit, students should be able to:
  • Distinguish between complete and reversible reactions. 
  • Explain dynamic equilibrium. 
  • State the characteristics of dynamic equilibrium. 
  • Explain the factors that affect the position of the equilibrium in a reversible 
  reaction. 
  • Apply Le Châtelier’s principle to explain the effects of changes in the 
  temperature, concentration and pressure on a system in equilibrium. 
  • Compare and contrast theoretical and actual optimal conditions in the 
  industrial processes. 
  • Relate the effect of concentration, temperature, pressure and catalyst to the 
  amount of products in the manufacturing industries. 
  • Recognize the importance of Le Châtelier’s principle in Haber and Contact 

  processes. 

  

  

  The above figure shows that when two teams pull on a rope with equal force. The 
  resulting force is equal in magnitude and equal to zero and the rope does not move, 
  the system is said to be in equilibrium. Students in figure (a) represent a system 
  in equilibrium. The equal and opposite forces on both ends of the seesaw are 
  balancing. If, instead one force is greater in magnitude than the other, the system 
  is not in equilibrium [ figure (b)]In chemistry, a chemical reaction is a process where old bonds are broken and new 
  bonds are formed. For a chemical reaction to take place, two or more substances 
  called reactants are interacted. In general, when reactants collide with sufficient 
  energy and in a proper orientation, the products are formed. Many chemical 
  reactions proceed to a certain extent and stop. In some cases, reactants combine 
  to form products and the products also start combining to give back the reactants. 
  When such opposing processes take place at equal rates, no reaction appears and it 

  is said that a state of equilibrium has reached.

  15.1. Difference between complete and incomplete reactions 
  (irreversible versus irreversible reactions)
  Activity 15.1
  1. Write any two equations of your choice to show a reaction that undergo completion.
  2. Write any two equations of your choice to show a reaction that does not go 
  completion
  A chemical reaction can proceed in either non-reversible (irreversible or 
  complete) or reversible reaction.
  During chemical processes, many chemical reactions do not undergo completion 
  but instead they attain a state of chemical equilibrium. Chemical reaction can 
  proceed in either non-reversible (irreversible or complete) or reversible reaction.
  A non-reversible reaction is a reaction which proceeds in only one direction, in 

  other words, the reactants are completely transformed into products.

  

  

  15.2. Concept of equilibrium (dynamic equilibrium) and its 
  characteristics 
  Activity 15.2
  1. Explain the terms used in equilibrium reactions.
  (a)Reversible reaction (b) equilibrium state (c) dynamic equilibrium (d) position of 
  equilibrium.
  2. Suggest and explain the characteristics of dynamic equilibrium and how it can 
  be attained.
  3. Learners should do a tug-of-war game outsidetheclassroom and comment on 
  the game.
  4. In a given Hotel, clients enter others leave. At a certain moment if the number 
  of leavers and arrivals is equal, the number of the clients in the Hotel doesn’t 
  change.
  i. Has the movements of clients coming in and out stopped?
  ii. How can you qualify that status?

  iii. How can you compare this with chemical equilibrium?

  15.2.1. Concept of equilibrium reactions
  When a chemical reaction takes place in a container which prevents the entry or 
  escape of any of the substances involved in the reaction, the quantities of these 
  components change as some are being consumed and others are being formed at 
  the same time.
  Chemical equilibrium is the state at which the rate of forward reaction becomes 
  

  equal to the rate of backward reaction.
  

  
  

  At the initial state, the rate of forward reaction is greater than the rate of backward 
  reaction. However as the products are formed, the concentration of reactants 

  decreases and the concentration of products increases.

  

  

  Figure 15.5 variation of concentration of A and B with time for reversible reaction 

  Consider the reaction A B; the figure15.4 (b) indicates how the concentration 
  of A decreases while that of B increases for the reaction. The dotted vertical line 
  indicates the time when the concentrations of A and B are no longer changing.
  If the reversible reaction is carried out in a closed system, the reaction is said to be in 
  the equilibrium state when the forward and backward reaction occur simultaneously 
  at the same rate and the concentrations of reactants and products do not change 
  with time (Figure 15.4 b). 
  At this point, the rates of forward and reverse reactions are the same and the system 
  is said to have reached a state of dynamic equilibrium.
  A dynamic equilibrium is a process where the forward and reverse reactions 
  proceed at the same rate;at that moment the concentrations of reactants and 
  

  products remain constant (do not change).

  However, in dynamic equilibrium, even if the concentrations of reactants and 
  

  
  

  products do not change, it does not mean that the reaction has stopped. Rather, the 
  reaction is proceeding in a way that it keeps the concentrations unchanged (the net 
  change is zero).
  There are two types of chemical equilibrium: homogeneous and heterogeneous
  equilibria.
  In a homogeneous equilibrium, all the reactants and the products are in the same 
  

  phase. 

  
  

  
  

  Checking up 15.2
  1. Briefly explain the characteristics of reactions at equilibrium
  2. Compare the homogeneous and heterogeneous reactions using specific 
  examples.
  3. By giving an example, describe the term dynamic equilibrium.
  4. When does a reaction attain equilibrium state?
  5. Using a graph and specific examples, explain what happens during a reaction 
  

  before, at and after the equilibrium has been attained. 

  Activity 15.3(a)
  

  
  

  1. Around 1908-1909 a young German research chemist, Fritz Haber, had discovered that nitrogen and hydrogen would form an equilibrium mixture containing ammonia.
  (a) Write a balanced equation for the formation of ammonia.
  (b)Haber’s experiment yielded an equilibrium mixture containing only 8% by volume of ammonia. What conditions of temperature and pressure does Le Châtelier’s principle predict for maximum yield of ammonia at equilibrium?
  (c) Why do you think Haber employed the catalyst accompanied with promoters 
  and heat exchanger in his equipment?
  2. How is Le Châtelier’s principle used to explain the conditions that affect the 
  equilibrium reactions?
  Many industrial processes involve reversible reactions. It is important to understand 
  how the variation of conditions can affect the composition of a chemical equilibrium. 
  Some reactions to take place involve some conditions. For example, the rate of a 
  chemical reaction depends on factors that affect the reaction. 
   Different factors which can affect the chemical equilibrium include:
  1. Temperature
  

  
  2. Pressure
  3. Concentration of reactants and products
  The effect of the above-mentioned factors on chemical equilibrium can be explained 
  by the Le Châtelier’s Principle.
  Le Châtelier’s Principle
  According to Le Châtelier’s Principle, when the temperature, pressure or concentration 
  of a reaction in equilibrium is changed, the reaction shifts in the direction where the 
  effect of these changes is reduced.
  15.3.1. Effect of Temperature on equilibrium
  Activity 15.3(b)
  1. Explain the following terms
  (a)Endothermic (b)Exothermic
  (c) Suggest how temperature affects the position of equilibrium.
  When dealing with temperature, we distinguish exothermic and endothermic 
  reactions. A change in the temperature of a system already in equilibrium could 
  either shift the equilibrium to the right (favoring the forward reaction) or to the left
  (favoring the backward reaction). This depends on whether the forward reaction is 
  exothermic or endothermic. Heat can be considered a reactant in an endothermic 
  reaction and a product in an exothermic reaction. For a reversible reaction, when 
  the forward reaction is exothermic, the enthalpy change is negative (ΔH < 0), then 
  the backward reaction is endothermic and the enthalpy change is positive (ΔH > 0).
  For exothermic forward reactions, an increase in temperature will cause the system 
  to counter balance it by favouring the reaction that consumes heat, hence the 
  backward reaction will be favoured or promoted. On the contrary, if the temperature 
  is decreased, the system reacts to produce more heat by favouring the forward 
  

  reaction.

  

  

  

  15.3.2. The effect of change in concentration on equilibrium
  Activity 15.3(c):
  Experiment to investigate the effect of changing concentration on 
  equilibrium
  Equipment/materials
  
  
  
  
  According to Le Châtelier’s principle, an increase in pressure favours the reaction in 
  the direction where the volume of reactants is reduced, or less molecules of gas are 
  formed, and a decrease in pressure favours the reaction in the direction where the 
  

  volume of reactants is increased, or more molecules of gas are formed.

  
  

  
  To obtain much ammonia in the equilibrium mixture, a high pressure of 200 
  atmospheres is needed.
  

  The effect of pressure can be summarized by the graph indicated below

  
  High Pressure gives a good yield of ammonia as indicated from the graph above, at 
  400 atmosphers the yield of ammonia is 70%
  Higher pressure increases the rate of reaction
  However, the higher the pressure used, the higher the cost of the equipment needed 
  to withstand the pressure.
  The higher the pressure the higher the electrical energy costs for pumps to produce 
  the pressure.
  A moderately high pressure of between 150 – 300 atmospheres is used. 
  
  15.3.4. The effect of a catalyst on equilibrium
  Activity 15.3.4
  1. What is an enzyme?
  2. What is a catalyst? Name the catalyst used in the Haber process and contact 
  process
  3. What would happen if the enzymes involved in the digestion of food were not 
  present?
  4. Most of the metabolic processes in the body are controlled by enzymes. What 
  would happen to these metabolic processes if the enzymes were missing?
  The function of a catalyst is to speed up the reaction by lowering the activation 
  energy. The catalyst lowers the activation energy of the forward reaction and reverse 
  reaction to the same extent. Adding a catalyst doesn’t affect the relative rates of the 
  two reactions and therefore the catalyst has no effect on the equilibrium system. 
  But the catalyst helps the system to reach the equilibrium more quickly. The catalyst 
  does not appear in the overall equation of the reaction.
  Practical and financial aspects: In industry, all the above factors must be considered, 
  taking in account not only the theoritical advantages but also their costs and risks. 
  

  That is why for example the manufacture of ammonia is based on a compromise of 

  

  

  

  15.4. End unit assessment

  
  

  
  

  
  

  
  

  
  

  
  

  
  

  
  

• UNIT 16:ACIDS AND BASES

  UNIT 16: ACIDS AND BASES
  Key unit Competence: Explain the acid-base theories (Arrhenius, Bronsted–Lowry, 
  Lewis).
  Learning Objectives: 
  By the end of this unit, students should be able to:
  • Explain the concept of acid and base using Arrhenius, Brønsted-Lowry and 
  Lewis’ theory. 
  • Distinguish strong acids from weak acids and strong bases from weak bases 
  using Brønsted-Lowry theory.
  • Classify the acids and bases as strong or weak according to their dissociation 
  in aqueous solution.
  • Distinguish between Brønsted-Lowry and Lewis’ Acid-Base theories.
  • Write the dissociation of acids and bases and identify the acid-base 
  

  conjugate pairs

  

  

  

  

  16.1.1. Arrhenius Theory of Acid-Base
  The first person to recognize the essential nature of acids and bases was the Swedish 
  

  scientist Svante Arrhenius (1859–1927). On the basis of his experiments with

  

  

  

  

  

  a. Two strongest acidic substances
  b. Two weakest acidic substances
  c. Two most alkaline substances
  d. Two least alkaline substances
  

  e. Neutral substance (s)

  

  

  

  

  

  

  

  

  

  16.2. End unit Assessment

  
  

  
  

  
  

• UNIT 17:REDUCTION AND OXIDATION REACTION

   UNIT17: REDUCTION AND OXIDATION REACTION
  Key unit competency: Explain the concept of reduction and oxidation and balance 
  equations for redox reactions
  Learning Objectives
  By the end of this unit, students should be able to
  • Explain the redox reactions in terms of electron transfer and changes in 
  oxidation state (number).
  • Explain the concept of disproportionation 
  • Differentiate the reducing agent from the oxidizing agent in a redox 
  reaction.
  • Work out the oxidation numbers of elements in the compounds. 
  • Perform simple displacement reactions to order elements in terms of 
  oxidizing or reducing ability. 
  • Apply half-reaction method to balance redox reactions. 
  

  • Deduce balanced equations for redox reactions from relevant half equations. 

  

   

  17.1. Definition of electrochemistry and its relationship with 
  redox reactions. 
  Activity 17.1. 
  1. Use examples to differentiate redox reactions from other chemical reactions
  2. Explain this statement: “Electrochemistry is a chapter of chemistry that studies 
  the chemical reactions that produce electricity”
  Electrochemistry is defined as the study of the interchange of chemical and Electrical 
  energy. It is primarily concerned with two processes that involve oxidation–reduction 
  reactions: the generation of an electric current from a spontaneous chemicalreaction 
  and, the opposite process, the use of a current to produce chemical change.
  Electrochemistry is important in other less obvious ways. For example, the corrosion 
  of iron, which has tremendous economic implications, is an electrochemical process. 
  In addition, many important industrial materials such as aluminum, chlorine, and 
  sodium hydroxide are prepared by electrolytic processes.
  
  

  

  

  

  

  Hence a redox reaction is a combination of two half-reactions: an oxidation halfreaction and a reduction half-reaction. Nevertheless, one half-reaction cannot exist 
  without the other, because electrons lost in the oxidation process must be captured 
  in the reduction process, this explains why we talk of oxidation-reduction or redox 
  reaction. 
  The characteristic of a redox reaction is that there is exchange or transfer of electrons 
  between chemical species participating in the reaction.
  We can compare this to the emigration-immigration movement: when a person 
  leaves a country, emigration for that country, he/she must enter another country, 
  immigration for that country and this constitutes an emigration-immigration 
  movement.
  

  We notice that any chemical species whose oxidation state increases is oxidized: 

  

  1. Aqueous copper (II) ion reacts with aqueous iodide ion to yield solid copper (I) 
  iodide and aqueous iodine. 
  a. Write the net ionic equation, 
  b. Assign oxidation numbers to all species present, and 
  

  c. Identify the oxidizing and reducing agents.

  

  

  

  

  The oxidation number of an atom is the apparent or real charge that the atom has 
  when all bonds between atoms of different elements are assumed to be ionic. By 
  comparing the oxidation number of an element or chemical species before and 
  after reaction, we can tell whether the atom has gained or lost electrons. Note that 
  oxidation numbers don’t necessarily imply ionic charges; they are just a convenient 
  

  device to help keep track of electrons during redox reactions.

  

  

  

  

  

  17.5. Balancing of redox equations

  
  

  3. To the second portion add a few drops of hydrogen peroxide followed by one 
  or two drops of dilute surphuric acid and warm gently. Allow the solution to 
  cool (or cool it under running tap water). To the cold solution add drop wise 2M 
  NaOH until there is no further change. Record your observations. Add dilute 
  

  sulphuric acid to the resultant product and note down your observations. Rinse 
  the test tube thoroughly
  4. To the third portion, add about 1 cm3
   of dilute hydrogen peroxide solution followed by one or two drops of dilute sulphuric acid. Warm gently and test the 
  gas produced with a glowing splint. Allow the solution to cool (or cool it using 
  running tap water).To the cold solution add ammonia solution drop wise until 
  no further change. Compare the product formed when ammonia solution to 
  

  that obtained when sodium hydroxide was used.

  Study Questions 
  1. Name the products formed when dilute sulphuric acid reacts with iron powder. 
  Write a balanced formula equation for the reaction
  2. When dilute sulphuric acid reacts with iron powder, iron atoms are oxidized and 
  hydrogen ions are reduced. Write a balanced 
  a) oxidation half-equation
  b) reduction half-equation and 
  c) overall redox equations for the reaction between iron and sulphuric acid
  3. What is the effect of adding a hydrogen peroxide in step 4? 
  4. What will be the effect of adding concentrated nitric acid to any iron salt? 
  Explain why concentrated nitric acid does not react with pure iron metal
  17.5.1. Rules for balancing redox reactions
  The Half-Reaction Method for Balancing Oxidation–Reduction Reactions in 
  Aqueous Solutions
  For oxidation–reduction reactions that occur in aqueous solution, it is useful to 
  separate the reaction into two half-reactions: one involving oxidation reaction and 
  the other involving reduction reaction. Then after balancing those half reactions, 
  find the overall oxidation-reduction (redox) reaction by combining the two halfreactions. 
  For example, consider the unbalanced equation for the oxidation– reduction 
  reaction between cerium(IV) ion and tin(II) ion:
  
  
  
  
  
  
  
  
  
  
  
  2. Place the test tubes in a 400 mL beaker that is about 1/3 full of boiling water. 
  After
  a few minutes, look for evidence of reaction. Note any changes. Did some metals 
  that didn’t react with cold water, react with hot water? 
  3. Place a small sample of each metal in test tubes containing 5 mL of 1.0 mol/L 
  hydrochloric acid, HCl. Watch for evidence of reaction. Note any changes
  4. Place a small sample of magnesium ribbon in test tube containing 5 mL of 1M 
  copper (II) sulphate. Watch for evidence of reaction and note any changes
  Study questions
  1) Considering sodium, magnesium, zinc, and copper: 
  Arrange the metals in order of increasing reactivity (from least reactive to most reactive)
  2) Which of the four metals are reacting with cold water? For those metals that did 
  react, write a balanced symbolic equation.
  3) Which of the four metals are reacting with hot water? For those metals that did 
  react, write a balanced symbolic equation. 
  4) Which of the four metals are reacting with the hydrochloric acid? For those 
  metals that did react, write a balanced symbolic equation. 
  5) Which metal did not react with either water or hydrochloric acid?
  6) Which of the four metals would be suitable for making saucepans? Explain why 
  the others are not.
  7. Describe what you would see if you dropped a piece of magnesium ribbon into 
  some copper (II) sulphate solution in a test tube. Write a chemical equation for 
  the reaction.
  The reactivity series is a series of metals, in order of reactivity, as reducing agents, 
  from highest to lowest reducing agent. It is used to determine the products of single 
  displacement reactions, whereby metal A will displace another metal B in a solution 
  if A is higher in the series. Although hydrogen is not a metal, it is included in the 
  reactivity series for comparison (Table 17.2).
  When a metal is placed in a solution of another metal salt, and if the metal is more 
  active than the metal in the salt, the more active metal displaces the other metal 
  
  
  
  
  17.7. End unit assessment
  
  
  
  
  
  
  
  b) Write balanced equations for each reaction that took place. 
  18. Sulfur dioxide reacts with water to form sulfite ion. Is this a redox reaction? 
  Justify your answer.
  19. In each of the following balanced redox equations, identify:
  i) the species oxidized and their new oxidation numbers
  (ii) the species reduced and their new oxidation numbers.
  (iii) the reducing agent
  

  (iv) the oxidizing agent

  

  i) copper + chromium sulfate 
  ii) magnesium + chromium sulfate 
  iii) chromium + copper sulfate 
  

  d. Compare the reactivity of chromium with those of iron and zinc

• UNIT 18:ENERGY CHANGES AND ENERGY PROFILE DIAGRAM

  UNIT 18: ENERGY CHANGES AND ENERGY PROFILE 
  

  DIAGRAMS
  

  Key unit Competence: Explain the concept of energy changes and energy profile 
  

  diagrams for the exothermic and endothermic processes.
  

  Learning Objectives 
  

  By the end of this unit, student should be able to:
  

  • Define the term Thermochemistry.
  

  • Explain the concept of system and distinguish between the types of systems.
  

  • Distinguish between Temperature and heat.
  

  • Explain the concept of Exothermic and endothermic reactions and represent 
  

  them using energy profile diagrams.
  

  • Carefully deal with reactions that produce a lot of energy.
  

  • Appreciate the use of chemical energy in daily life.
  

  • Respect the experimental protocol during chemistry practicals.
  

  • Relate the type of reaction to its energy profile diagram.
  

  • Interprete the experimental results about energy changes occurring during 
  

  chemical reactions.
  

  • Explain the energy change as a function of the breaking and formation of 
  chemical bonds.
  
  
  
  18.1. Concept of a system
  18.1. Concept of a system
  

  Activity 18.1
  

  Topic: Energy transfer between a system and surroundings.
  

  Apparatus and equipment (per group)
  

  • Eye protection
  

  • Four test-tubes or four expanded polystyrene cups with lids to act as 
  

  calorimeters
  

  • Spatula
  

  • Teat pipette or small measuring cylinder
  

  • Thermometer
  • Access to a balance
  
  

  Experiment 4
  

  Repeat the same procedure as in experiment 1 but use copper (II) sulphate 
  

  solution and Zinc powder instead of water and anhydrous copper (II) sulphate respectively.
  

  Safety
  

  • Wear eye protection.
  

  • Anhydrous copper (II) sulfate is harmful.
  

  • Zinc powder is flammable.
  

  Introduction
  

  Instant hot and cold packs are available for use in first aid. This experiment 
  illustrates the types of chemical reaction that occur in these packs.
  
  
  

  Study questions
  

  1. Identify the reactions that are exothermic and those that are endothermic.
  

  2. Write symbol equations to represent the chemical reaction taking place in Experiment 3.
  

  3. Which two substances could be put in a cold pack?
  

  4. Golfers need a hand warmer to keep their hands warm on a cold day. Which 
  

  chemicals could be put in these warmers?
  

  All chemical reactions involve the breaking of bonds in the reactants and the 
  

  formation of new bonds in the products. The breaking of bond requires energy, 
  

  whereas the formation of bond releases energy.
  

  Thermochemistry is the study of heat and energy associated with a chemical 
  

  reaction or a physical transformation. Thermodynamics is the study of the 
  

  relationship between heat, work, and other forms of energy. A reaction may release 
  

  or absorb energy, and a phase change may do the same, such as in melting and 
  

  boiling. Energy is exchanged between a closed system and its surroundings during 
  

  the heating and cooling processes.
  

  A system is a part of the universe which is studied using laws of thermodynamics. 
  

  Everything outside the system is the surroundings. An infinitely small region 
  

  separating the system from the surroundings is called boundary. In Chemistry the 
  

  chemical system consists of reactants and products. The systems are classified 
  

  according to the number of factors including the composition and the interaction 
  

  with the surroundings. A system can be homogeneous or heterogeneous. It can 
  

  be in gaseous, liquid or solid state. A system is said to be in equilibrium when its 
  

  properties do not change with time. The state of a system is described using its 
  

  composition, temperature and pressure.
  

  Three types of systems can be distinguished according to the exchange between 
  

  the system and the surroundings in terms of matter and/ or energy. 
  

  1. An open system is a system that can exchange both matter and energy with the 
  surroundings (Figure 18.1).
  
  
  

  lid prevents the exchange of matter between the system and the surroundings. An 
  

  isolated system is a system which is both sealed and insulated. It can exchange 
  

  neither matter nor energy with its surroundings. 
  

  Examples
  

  Hot coffee in a thermos flask (Figure 18.3).The latter is a closed system. The outer 
  

  surface is insulated and thus neither heat nor matter transfer take place between 
  the system and the surrounding.
  
  
  

  2. Indicate the direction of heat (from one compartment to another) and explain 
  

  your answer for the following phenomenon
  

  a) When you touch water in a saucepan on top of a stove with your hand and you 
  

  fill it is warm
  

  b) When you touch water from the tap with your hand and you fill it is cold
  

  c) When you mix cold water and warm water
  

  18.2.1. Internal energy
  

  The first Law of Thermodynamics deals with energy that is transferred between a 
  

  given system and its surroundings in form of heat. The exchange of energy is related 
  

  to the energy that is stored in the system called internal energy E. The internal 
  

  energy is the sum of the kinetic and potential energies of the particles that form a 
  system.
  
  

  18.2.2. Heat energy and temperature
  

  The heat or thermal energy of an object is the total energy of all the molecular 
  

  motion inside that object. When two bodies are in contact, heat always flows from 
  

  the object with the higher temperature to that of lower temperature. Heat transfer 
  

  ceases when a thermal equilibrium is attained. The heat content of a body will 
  

  depend on its temperature, its mass, and the material it is made of. Because heat is a 
  

  form of energy, it is measured in Joules (J) or kilojoules (kJ) or calorie (cal). A calorie 
  

  is defined as the amount of energy needed to raise the temperature of one gram of 
  

  water by one degree Celsius.
  

  1 calorie (cal) = 4186 joules (J); 1000 cal = 1 kcal = 4.186 kJ.
  

  The temperature is a measure of the average heat energy (thermal energy) of the 
  

  molecules in a substance. When an object has a temperature of 100 °C, for example, 
  

  it does not mean that every single molecule has that exact thermal energy. In any 
  

  substance, molecules are moving with a range of energies, and interacting with 
  

  each other. The temperature is a physical measure expressing how an object is hot 
  

  or cold.The temperature is measured using a variety of temperature scales. The most 
  commonly used are degree Celsius (°C) and degree Kelvin (K):
  K = °C + 273
  

  N.B: In thermodynamic calculation, degree Kelvin, not degree celcius, is used.
  

  First Law of Thermodynamics
  

  Thermodynamics  is part of physical chemistry that deals with the relationships 
  

  between heat and other forms of energy. In particular, it describes how thermal 
  

  energy is converted to and from other forms of energy and how it affects matter. 
  

  The first Law of thermodynamics is a statement about conservation of energy and it 
  

  categorizes the method of energy transfer into two basic forms: work (W) and heat 
  

  (Q). The First Law of Thermodynamics states that energy can be converted from 
  

  one form to another with the interaction of heat, work and internal energy, 
  

  but it cannot be created or destroyed, under any circumstances. Internal energy 
  

  refers to all the energies within a given system, including the kinetic energy of 
  

  molecules and the energy stored in all of the chemical bonds between molecules. 
  

  For a closed system (without mass input and output), the internal energy is the sum 
  

  of the heat energy and the work done by the system or the surroundings 
  ∆U = Q + W
  

  Where W is the energy transferred to the system by doing work and Q is the energy 
  transferred to it by heating. 
  
  
  

  The work done by the system on the surroundings is negative. Therefore, the first 
  

  law of Thermodynamics is written as:
  

  ΔU = Q – W
  

  Work (W) is also equal to the negative external pressure on the system multiplied by 
  

  the change in volume. It can be expressed as: 
  

  W = −P∆V
  

  Where P is the external pressure on the system, and ΔV is the change in volume. 
  

  This is specifically called pressure-volume work. Therefore, the Fist Law of 
  

  Thermodynamics is expressed using equation: 
  ΔU = Q -P∆V
  
  
  
  

  Glasses P and Q have the same amount of water. Glasses R and S have the same 
  

  amount of water.
  

  The water in Glasses P and R are at the same temperature. The water in Glasses Q 
  

  and S are at the same temperature. 
  

  1. Fill in the blanks below with the correct answers.
  

  a. The water in Glass……..has the most heat. 
  

  b. The water in Glass……..has the least heat.
  

  2. Ari touched a metal spoon. The metal spoon felt cold. Choose the best answer.
  

  a. Heat flows from hand to spoon
  

  b. Heat flows from spoon to hand
  

  c. Heat does not flow
  

  d. Heat flows in both directions
  

  3. Tom placed a metal spoon in a mug of hot coffee as shown below. The metal 
  spoon got hot. Choose the best answer
  
  
  
  

  a. Heat flows from hand to spoon
  

  b. Heat flows from spoon to hand
  

  c. Heat does not flow
  

  d. Heat flows in both directions
  

  4. Complete the statement below.
  

  If two objects are near each other and one object is hotter than the other, then 
  

  heat will flow from the …………………….object to the…………………..
  

  object.
  5. Complete the crossword puzzle using the clues given below.
  
  

  Down
  

  1. Our sense of ………………….cannot measure temperature accurately.
  

  3. Wood is a …………………….conductor of heat.
  

  4. Heat is a form of ………………………..
  

  6. ………………………….is a measure of how hot or cold an object is.
  

  10. Metals can ………………………………when heated.
  

  Across
  

  2. Heat is used to …………………………… food.
  5. When two objects of different temperatures are in contact, heat will travel from 
  the ………… object to the other object.
  6. What does the first law of thermodynamics have to do with systems?
  7. The instrument used to measure temperature accurately is a 
  ……………………………..
  8. Temperature is measured in the unit ……………………….Celsius (°C).
  9. A……………, when used with a temperature sensor, can be used to measure 
  and record temperatures.
  10. The Sun is an important ………………………….of heat.
  11. A hotter object will has a ……………………….temperature.
  12. A gas is compressed and during this process the surroundings does 462 J of 
  work on the gas. At the same time, the gas loses 128 J of energy to the surroundings 
  as heat. What is the change in the internal energy of the gas?
  

  18.3. Standard Enthalpy changes 

  Activity 18.3 
  1. What is meant by standard conditions of temperature and pressure?
  

  2. Which term describes the sum of kinetic energy and potential energy?

  

  

  

  

  

  

  

  The standard enthalpy of Hydration also called Standard enthalpy of solvation
  is the amount of heat released when one mole of isolated gaseous ions dissolve in 
  water forming one mole of aqueous ions under standard conditions. The positive 
  terminal of the water molecule is attracted to the anion while its negative terminal 
  is attracted to the cation. This is an ion-dipolar attraction which is typically an 
  

  electrostatic interaction. This latter is accompanied by the release of heat energy.

  

  

  

  

  1. Discuss the type of energy form present in points A, B and C of the pathway 
  followed by the vehicle.
  2. Discuss how each form of energy changes from point A to point C.
  3. Which points corresponds to maximum stability and minimum stability, 
  respectively? Relate your answer to energy concept.
  When a chemical reaction happens, the energy is transferred to or from the 
  surroundings and often there is a temperature change. For example, when a bonfire 
  burns, it transfers the heat energy to the surroundings. The objects near the bonfire 
  become warmer and the temperature rise can be measured with a thermometer.
  There are some chemical reactions that must absorb energy in order to proceed. 
  These are endothermic reactions. Some other chemical reactions release energy to 
  the surroundings. The energy released can take the form of heat, light, or sound. 
  These are exothermic reactions.
  1. Exothermic reactions
  They are characterized by an increase in the temperature of the surroundings, i.e. 
  energy is given up. Heat is lost to the surroundings and by convention it is negative 
  and represented as: ΔH < 0
  For exothermic reaction (Figure 18.8), total energy of the reactants is higher than in 
  the product, because the heat energy absorbed during bond breaking is lower than 
  

  the heat energy released during bond formation.

  

  

  2. Endothermic reactions
  These are reactions that take place by absorbing the energy from the 
  surroundings. The energy is usually transferred as heat energy; in this case the 
  surroundings loses energy to the reactants causing the surroundings to get colder.
  Endothermic reactions cannot occur spontaneously. Work must be done in order 
  to get these reactions to occur. When endothermic reactions absorb energy, a 
  temperature drop in the surroundings is observed during the reaction. Endothermic 
  reactions are characterized by positive heat flow (into the reaction) and an increase 
  in enthalpy, by convention it is represented by: ΔH > 0
  For endothermic reaction (Figure 18.9), the total energy of the reactants is  lower 
  than the product, because the heat energy absorbed during bond breaking is higher 
  than the heat energy released during bond formation.
  You have certainly experienced this effect when you put a drop of methanol or any 
  other volatile substance on your skin; you feel cold because that part of your skin is 
  

  supplying energy to evaporate the volatile liquid

  

  

  3. Activation energy, Ea
  The activation energy is the minimum energy required for a chemical reaction to 
  take place. It is the energy barrier that has to be overcome for a reaction to proceed. 
  Without that minimum energy, the reaction will not take place. That is why, for 
  example, the only fact that a dry wood is in contact with oxygen of air will not 
  start burning; there is a need of supplying the minimum energy to overcome the 
  activation energy barrier, this is done by using a burning match.
  4. Activated complex
  The activated complex is the intermedicate species, where former chemical bonds 
  are being broken, whereas new chemical bonds are being formed. In term of energy, 
  

  it corresponds to the activation energy.

  

  

  

  

  

  5. Determine the activation energy for the reverse reaction.
  6. Determine the enthalpy change of reaction for the forward reaction.
  7. Determine the enthalpy change of reaction for the reverse reaction.
  8. Fill in using exothermic or endothermic.
  a. The forward reaction is ……………………..
  b. The reverse reaction is ………………………
  9. Which chemical species or set of chemical species represent the activated 
  complex?
  10. Which one of the chemical bonds A-X and M-X is stronger? Explain.
  11. State the chemical species whose particles move the fastest. Explain your 
  answer.
  12. State the chemical species whose particles move the slowest. Explain your 
  answer.
  13. The compound AX and the element M are in gaseous and solid states, 
  respectively. 
  What effect would grinding M into a fine powder have on this energy profile 
  

  diagram?

  18.5. End unit Assessment

  
  

  Regarding the absorption or release of energy, what is the nature of the 
  overallreaction?
  b. What is the activation energy for the forward reaction?
  c. What is the activation energy for the reverse reaction?
  d. Determine the enthalpy change of reaction for the forward reaction?
  e. Is the reverse reaction endothermic or exothermic?
  f. Which chemical species constitute the activated complex?
  g. Which chemical species or set of chemical species have the maximum potential
  energy?
  h. Which chemical species or set of chemical species have the maximum kinetic 
  

  energy?

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• Topic 19