Chemistry

Key unit competency

Use atomic structure and electronic configuration to explain the trends in the physical properties of elements.

Learning objectives

By the end of this unit, I will 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).

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

149ChemistrySenior Four Student's Book5.1. Historical Background of the Periodic TableActivity 5.1Who is the father of the periodic table? Explain your answers.Differentiate the laws of triads and octavesDuring 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.

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.

n 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.

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 1680Robert 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 1932James 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.

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

Theperiodic 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 elements 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.

a) Finding the Period of elements

The Period of an element is equal to the highest energy level of electrons or principal quantum number which is being filled by electrons.

For a better understanding the following is an example:

Consider the element 16S with the electronic configuration: 1s22s22p63s23p4, the number 3 is the highest energy level or principal quantum number of electrons. Thus the period of S is 3.

If we consider another example, the electronic configuration of 23Cr is: 1s22s22p63s23p64s13d5 The number 4 is the highest energy level or principal quantum number of electrons. Thus the period of Cr is 4.

b) Finding the Group of elements

The Group of an element is equal to the number of outermost or valence electrons of element or number of electrons in the highest energy level of elements. Another way of finding the group of element is looking at sub shells. If the last sub shell of electron configuration is “s” or “p”, then the group becomes “1 to 18” with the number of group corresponding to the electrons occupying the last orbitals according to the given examples.

Consider the element 19K with the electronic configuration:1s22s22p63s23p64s1.Since the last sub shell is “s” with one electron, the element K is located in group 1of the periodic table.

Similarly, the element 35Br has the corresponding electronic configuration:1s22s22p63s23p64s23d104p5. Since the last sub shell is “p”, the element bromine (Br) is located in group 17of the modern periodic table.

The elements of groups 3 to 12 have the electron configuration ns and (n-1)d, total number of electrons in these orbitals. This informs us that those elements are located in group of element corresponding to transition metals.

Let us consider the following examples. The element 26Fe has the electronic configuration:1s22s22p63s23p64s23d6, 6+2 = 8 correspond to the group 8 of the transition elements.

Here are some evidences for everyone to find group number of elements.

Last Orbital Group: Considering the last orbital of an element, we are able to identify in which group of the periodic table it is located. The following tables show the location of representative elements in different groups (Table 5.2) and transition elements (Table 5.3 based on the last orbital.

Table 5.2. Group number of representative elements

For the transistion elements or d block elements, their valence electron configuration is ns1-2and (n-1)d1-10, hence the total number of valence electrons in these orbitals gives the group number.

Example: Sc (4s23d1), its group number is 3

Cu (4s13d10), its group number is 11

Table 5.3. Group number of transition elements

Example: Find the period and the group of the elements 16X and 24X using its electronic configuration:

16X: 1s22s22p63s23p4; the period of X is 3 and its group is 16

24X: 1s22s22p63s23p64s23d4; the period is 4 and the group is 4 + 2 = 6

The elements in the periodic table are also located using atomic orbitals.

Checking up 5.3

The elements X, Y, Z, T and U are given in the picture below.

Which one(s) of the following statements are correct, which one(s) are false for these elements.

a. X is alkaline metal

b. Y is in p block

c. Z is halogens

d. U is lanthanide

e. T is noble gas

2. Explain the major parts of the periodic table?

3. Differentiate the blocks included in a periodic table?

a. By using examples, locate at least two elements in each block

b. Indicate the group and the period of the chosen element.

5.4. Classification of Elements into Blocks (s, p, d, f-block)

Activity 5.4.

1. Apart from blocks, the elements in the periodic table are classified as representative elements, transition metals, lanthanides and actinides. Match these categories with s, p, d and f-blocks of the periodic table.

2. Classify the following elements into s, p, d or f-block and justify your answer: Al (z=13), K (z= 19), Ca (z =20) and Fe (z=26).

In the Periodic Table, the elements are organized into different blocks according to their electron configurations. They are classified into four blocks such as s-block, p-block, d-block and f-block depending on the type of atomic orbitals that are being filled with electrons. This division is based upon the name of the orbitals which 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 transitionmetals.

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.

5.5. Characteristics of different blocks of the periodic table

a. The s-block elements

The s-block comprises elements of Group 1 (alkali metals) and Group 2 (alkaline earth metals) which have ns1 and ns2 outermost electronic configuration. The elements of group 1 and 2 are classified as reactive metals with low ionization enthalpies and highly electropositive. They lose their valence electron(s) willingly to form a positive ion with the charge +1 in the case of alkali metals or +2 in the case of alkaline earth metals. The metallic character and reactivity of s-block elements increases as we move down the group. Because of high reactivity they are never found pure in nature. Most of the metals of this block give a characteristic flame color. They are soft metals with low melting and boiling points. The compounds of s-block elements, with the exception of those of lithium and beryllium are predominantly ionic. There are 14 s-block elements in the periodic table. The Table 5.4 below shows an example of the electronic configuration of some elements of Group 1.

Table 5.4. Electronic configuration ofthe group 1 elements

b. The p-block elements

The p-Block comprises elements belonging to Group 13 up to 18. The elements of p-block and s-block are called the representative elements or main group elements.

The valence electron of p-block elements varies from ns2np1 to ns2np6 in each period. At the end of each period appears a noble gas element with a closed valence shell ns2np6 configuration. The p-block includes 6 groups:13, 14, 15, 16, 17 and 18. The atoms of the elements belonging to these groups accept the last electron in 2p, 3p, 4p, 5p and 6p orbitals. Group 18 elements are called noble gases and all these elements have closed shell ns2 np6 electronic configuration in the outermost shell. All the orbitals in the valence shell of the noble gases are completely filled by electrons and it is very difficult to alter this stable arrangement by the addition or removal of electrons. The noble gases thus exhibit very low chemical reactivity.

The p-block elements consist of both metals and non-metals but the number of non-metals is considerably higher than that of metals. The metallic character increases as we move down the group and non-metallic character increases from left to right along a period. The ionization energiesof p-block elements are relatively higher compared to those of s-block elements. Most of them form covalent compounds. Some of p-block elements exhibit more than one oxidation state in their compounds and their oxidizing character increases from left to right in a period and reducing character increases from top to bottom in a group. Most of the p-block elements are highly electronegative and form acidic oxides. The p-block contains some elements that show intermediate properties between metals and non-metals; those elements are called “metalloid” or “semi-metals”, e.g. Si and Ge; they are also called “semi-conductors”.

c) The d-block elements

Elements in which the last electron enters any one of the five d orbitals are called d- block element. These elements comprise Group 3 to 12 in the centre of the Periodic Table (Figure 5.2). These elements are characterized by the filling of inner d orbitals by electrons and are therefore referred to as d-Block Elements. Their outer electronic configuration is (n-1)d1-10ns0-2. As these elements are transition metals, mostly of them form coloured ions and they display variable valences or oxidation states. However, Zn, Cd and Hg which have the electronic configuration, (n-1) d10ns2 do not show most of the properties of transition elements.

The properties of d-block elements are midway between those of s- block and p- block elements, they are divided into four series called 1st, 2nd, 3rd and 4th transition series. The first transition series contains 10 elements from scandium (Z=21) to zinc (Z=30); for those elements, 3d orbitals are being progressively filled in. The second transition series contains the elements from Y through Ag in which the 4d orbitals are being progressively filled. The third transition series also contains the elements La and the element Hf through Au in which the 5d orbitals are progressively filled. The fourth transition series comprise elements also encloses the elements from Rf to Rg with 6d orbitals being filled.

The elements of d-block are hard, malleable and ductile metals with high melting and boiling point. Mercury is an exception, it is liquid at room temperature. Because they are metals, they are good conductors of heat and electricity.

Their ionization energy are between s and p block elements. Transition metals show variable oxidation states reason why they form both ionic and covalent compounds. In chemistry, catalysts used in different reactions are transition metals (i.e: V, Cr, Mn, Fe, Co, Ni, and Cu).

d. The f-Block elements

The f-block elements are found within the two rows of elements at the bottom of the Periodic Table. They are called the Lanthanides and include Ce (Z = 58) to Lu (Z = 71) and Actinides, Th (Z = 90) to Lr (Z = 103). All of these elements are characterized by the outer electronic configuration (n-2)f1-14 (n-1)d 0–1 ns2. The last electron added to each element is filled in f- orbital. These two series of elements are hence named the inner transition elements (f-block elements). They are all metals and display similar properties within each series. The f-block elements exhibit variable oxidation states, they have high melting and boiling point. They form complex compounds and most of the elements of the actinide series are radioactive.

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, electro positivity, 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 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 affect their ability to attract electrons easily as for 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 fluorineand 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.

Table 5.5. First ionization energies (kJ/mol) for the alkali metals and noble gases

As mentioned above, different types of ionization energy exist such as first and second ionization energy. In general, the second ionisation energy of an element is always greater than the first ionisation. This is explained as follows: every time you remove an electron from an atom, the remaining electrons are more strongly attracted by the nucleus and it require more energy to remove other electrons from the atom.

Hence: 1st IE < 2nd IE < 3rd IE

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 non-metallic 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. Elctrical 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 lose 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)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.