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 holdatoms 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 inthe 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 thatcontains 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 calciumwith 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 oppositelycharged 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 isshown 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 theirmelting 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 belowCaution: 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 roomtemperature
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 solution5. 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 regulararrangement 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 anionic 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 ionsin 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 andglass
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 shouldconsider 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 metals3. 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 bestconductors 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 thatthe 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 continuesto 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 decreasesdown 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 ofdifferent 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 andnegative 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 solidsd. 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 above21. 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 dissociationenergies.) compare the values and give a detailed explanation.