CHAPTER 4.0 CHEMICAL BONDING

TUTORIAL 4.0
CHEMICAL BONDING

At the end of this topic, students should be able to:

4.1 Lewis structure a) State the octet rule.

b) Describe how atoms achieve stability by attaining stable

configuration of:
i. noble gas

ii. pseudo-noble gas; and

iii. half-filed orbital.

c) Describe the formation of the following bonds using Lewis dot
symbol.

i. ionic or electrovalent bond

ii. covalent bond

iii. dative or coordinate bond
d) Draw Lewis structure of molecules and polyatomic ions with

single, double and triple bonds.

e) Compare the bond length between single, double and triple

bonds.
f) Determine the formal charge and the most plausible Lewis

structure.

g) Explain the exception to the octet rule:

i. incomplete octet;
ii. expanded octet; and

iii. odd number electrons.

h) Illustrate the concept of resonance using appropriate examples.

4.2 Molecular shape and a) Explain Valence Shell Electron Pair Repulsion Theory (VSEPR).
polarity b) Draw the basic molecular shapes:

i. linear;
ii. trigonal planar;
iii. tetrahedral;
iv. trigonal bipyramidal; and
v. octahedral.
c) Predict the shapes of molecule and bond angles in a given
species.
d) Explain bond polarity and dipole moment.
e) Deduce the polarity of molecules based on the shapes and the

resultant dipole moment.

4.3 Orbital overlap and a) Illustrate the formation of sigma (σ) and pi (π) bonds
hybridisation from overlapping of orbitals.

b) Describe the formation of hybrid orbitals of a central atom:
sp, sp2, sp3, sp3d, sp3d2 using appropriate examples.

c) Illustrate the hybridisation of the central atom and the overlapping
of orbitals in molecules.

4.4 Intermolecular forces a) Describe intermolecular forces
i. van der Waals forces
• dipole-dipole interactions or permanent dipole
• London forces or dispersion forces

ii. hydrogen bonding.

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4.5 Metallic Bond b) Explain factors that influence van der Waals forces and hydrogen
bond.

c) Relate the effects of hydrogen bonding on following physical
properties:
i. boiling point
ii. solubility
iii. density of water compared to ice

a) Explain the formation of metallic bond by using electron sea
model.

b) Relate metallic bond to the properties of metal:
i. malleability;
ii. ductility;
iii. electrical conductivity; and
iv. thermal conductivity

c) Explain the factors that affect the strength of metallic bond
d) Relate boiling/ melting point to the molecular structure, types of

bonding and intermolecular forces for the elements of:
i. Period 3;
ii. Group 1; and
iii. Group 17.

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HOUR: 1

1. Table below shows the selected elements and their proton numbers:

Elements Proton Number
Al 13
P 15
Cl 17
Fe 26
Zn 30

a) Write the Lewis dot symbol for Al, P and Cl.

b) Write the electronic configuration for the cations formed by ALL elements.

c) State the type of stability of the cations formed from these elements. (CLO3, C3)

2. How would an atom acquire octet configuration when it forms bond with another atom?
PSPM JAN 2000 (CLO1, C2)

3. a) How is an ionic bond formed? (CLO3, C3)
b) By using Lewis dot symbol, show the formation of:
i) MgBr2
ii) Al2O3
iii) BaO

4. a) How is a covalent bond formed?

b) Use Lewis dot symbols to show the sharing of electrons between hydrogen and fluorine
atoms. Label the electron pairs as bonding pair(s) and lone pair(s). (CLO3, C3)

5. Describe the formation of dative bond by using NH4+, H3O+, Al2Cl6, and BF3 NH3, as examples.
(CLO3, C3)

6. The TABLE 1 below shows two elements with their respective proton numbers.

Elements U V
Proton Number 20 9

TABLE 1

Based on the table above, state the type of bond formed between element U and V. Explain how the

bond is formed. PSPM 2000 (CLO3, C3)

7. Elements X and Y have proton number of 12 and 7, respectively. Write the chemical formula of the
compound formed between X and Y. Show the formation of the compound using Lewis dot symbol.
(CLO3, C4)

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HOUR: 2

8. Draw the simplest Lewis structure for each of the following molecules. Determine whether these

molecules satisfy the octet rule. (CLO3, C4)
a) NCl3
b) ClF3
c) BeCl2
d) H2S
e) AlBr3
f) SeF4
g) NO

h) CN-
i) NO+

j) O3

9. Two resonance structures for NCO- are as follows:

Structure P Structure Q

N CO N CO

a) Determine the formal charge of each atom in both structures P and Q. (CLO3, C4)
b) Determine the most plausible structure. Explain.

10. Formaldehyde (CH2O), a liquid with an unpleasant odour with the skeletal structure

O

HCH

Draw the Lewis structure for formaldehyde and determine the formal charge of each atom.
(CLO3, C3)

11. Draw all possible Lewis structures for POCl3. Select the best Lewis structure. (CLO3, C3)

12. a) Define resonance structures. (CLO3, C4)
b) Draw the resonance structures for NO3− and SO3.

13. Dinitrogen oxide molecule, N2O, has three resonance structures. The atomic arrangement in the
molecule is NNO. One of the resonance structures is shown below:

NN O

a) Draw the other two resonance structures.

b) Calculate the formal charge of each atom in the resonance structures in 13. a)
c) Determine the most plausible Lewis structure between the two resonance structures. Give

your reasons. UPS 2012 / 2013 (CLO3, C4)

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HOUR: 3

14. Using VSEPR model, predict the molecular geometries of: (CLO3, C3)
a) O3
b) PBr5
c) NO2+
d) SnCl3-.

15. Predict the shape of SiF4, SF4 and XeF4 molecules. Explain why the shapes differ. (CLO3, C4)

16. Antimony pentafluoride, SbF5, reacts with XeF4 and XeF6 to form the ions XeF3+, SbF6- and XeF5+.
Describe the geometries of these ions. (CLO3, C3)

17. The nitronium ion, NO2+ is linear in shape but the nitrite ion NO2- is not linear. Explain the observation.
(CLO3, C3)

18. Iodine combines with chlorine to form ICI2+ and ICI2− ions.
a) Draw the Lewis structures and predict the shapes of both ions based on the Valence Shell
Electron-Pair Repulsion (VSEPR) theory.
b) Determine whether both the ions are polar or non polar. PSPM 2001 / 2002 (CLO3, C4)

HOUR: 4

19. For each of the molecules given; PCl3, PCl5 and POCl3,
a) Draw a Lewis structure and name the shape based on molecular geometry.

b) Predict the bond angle and deduce the polarity. PSPM 2012 /2013 (CLO3, C4)

20. a) Define dipole moment. (CLO3, C4)
b) How can a molecule be non-polar although it has polar bonds?
c) Explain whether each of the following molecules is polar or non-polar.
i) SO2
ii) HBr
iii) BF3
iv) SF6
v) NH3
vi) H2Se

21. An atom X has 5 valence electrons. X reacts with fluorine gas to form XF3 and XF5 compounds. For
each compound,

a) draw the Lewis structure,
b) predict the electron pair geometry and molecular geometry and draw the molecular geometry

and state the bond angle(s). PSPM 2015 / 2016 (CLO3, C4)

22. Xenon, Xe which is located in the noble gas group attains octet configuration. Xe forms compounds
such as the fluoride compounds, XeF2 and XeF4.
a) State the number of bonding pairs and lone pairs surrounding the central atom Xe in XeF2
and XeF4. How does Xe atom accommodate the total number of electrons.
b) Give the shapes of XeF2 and XeF4 molecules.
c) Besides XeF2 and XeF4, give a molecular formula of another fluoride of Xe.
PSPM JAN 2000 (CLO3, C3)

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HOUR: 5

23. a) Describe the formation of sigma and pi bonds. (CLO1, C2)
b)
By using valence bond theory, explain the bonding in HCl and show the bond formation by
the overlapping of atomic orbitals. (CLO3, C3)

24. a) What is meant by hybridisation? (CLO1, C1)
b)
Determine the hybridisation of the underlined atom in each of the following compounds.
i) BF3
ii) CCl4
iii) CH3CH=CH2 (CLO3, C3)

25. Describe the hybridisation process and draw the overlapping of orbitals for the formation of bonds

in each of the following molecules.

a) CH3F
b) AsCl5
c) XeF4
d) ICl3
e) BrF5
f) SeF4
g) HCN

h) OCN-

i) PH4+
j) ICl2– (CLO3, C4)

HOUR: 6

26. Based on the skeletal structure of bicarbonate ion, HOCOO–, draw and label the overlapping of
orbitals showing all the σ and π bonds in the anion. Estimate the O-C-O and C-O-H bond angles.
(CLO3, C4)

27. Cyanide ion, CN– is formed when the carbon atom undergoes hybridisation and bonds to nitrogen
atom. Show the hybridisation in CN– and draw the overlapping of orbitals.
PSPM 2006 / 2007 (CLO3, C4)

28. Formamide, I, is a clear liquid which is miscible with water and is normally used as a chemical
feedstock in pharmaceutical industries.

O

HC

NH2

Formamide, I

Draw the overlapping of hybrid orbitals for the bonding in I. Label all the σ and π bonds.
PSPM DK015 2014 / 2015 (CLO3, C4)

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HOUR: 7

29. a) What are dipole−dipole interaction? (CLO3, C3)
b) Explain how London forces arise between non-polar molecules. (CLO1, C2)
c) Give the factors that influence the strength of van der Waals forces.

30. Name the type of intermolecular forces that exist between molecules for each of the following

species:

a) I2
b) SO2
c) CH4
d) H2S
e) C6H6
f) H2O (CLO3, C3)

31. Which molecule has a higher boiling point? Explain. (CLO3, C3)
a) Br2 or ICl
b) CH4 or SiH4
c) HCl or Cl2

32. a) What is meant by hydrogen bond? (CLO1, C1)
b) Which has a higher boiling point, ammonia, NH3 or phosphine, PH3? (CLO3, C3)
c)
Arrange the following compounds in ascending order of boiling point. Explain. (CLO3, C3)

C2H5OH CH3OCH3 CH3COOH

HOUR: 8

33. The boiling points and solubility of ammonia, hydrogen chloride and fluorine are given below:

Compound Boiling point / oC Solubility in water / mol dm-3 at 298 K
hydrogen chloride −85 23
−188 0
fluorine −33 18
ammonia

a) Explain in terms of intermolecular forces the difference in boiling points between
i) hydrogen chloride and fluorine.

ii) ammonia and hydrogen chloride (CLO3, C3)

b) Explain why the solubility of ammonia and hydrogen chloride are so much higher than that
of fluorine. (CLO3, C3)

34. Explain why ice floats on water. (CLO3, C3)

35. Explain the following observation:

Methane, CH4 (16 g mol−1) has lower boiling point than propane, C3H8 (44 g mol−1) whereas water,
H2O (18 g mol−1) has higher boiling point than hydrogen sulphide, H2S (34 g mol−1).

PSPM 2015/2016 (CLO3, C3)

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HOUR: 9

36. a) Describe the metallic bond in aluminium by using electron sea model. (CLO3, C3)
b) Explain why the boiling point of calcium is higher than that of potassium. (CLO3, C3)

37. Using Band Theory, explain the electrical conductivity of Mg, Si and S. (CLO3, C3)

38. Arrange in order of increasing melting points for the solid crystals of copper, Cu, iodine, I2 and
diamond, C. Explain your answer with reference to the attractive forces between the atoms. Which
of the solid crystals can conduct electricity in the solid state? Explain.
PSPM JAN 2000 (CLO3, C4)

39. TABLE 2 shows the boiling point of the third period elements in the Periodic Table.

Element Na Mg Al Si P S Cl Ar

Boiling point 890 1120 2450 2680 280 445 -34 -186
(oC) TABLE 2
Based on the structures and bondings, explain the difference in the boiling points of the elements?

PSPM JAN 1999 (CLO3,C4)

40. Explain each of the following observations:
a) Boiling point of Cl2 (-34.6 ºC) is lower than of Br2 (58.8 ºC). PSPM 2006 / 2007 (CLO3, C4)
b) The melting point of sodium is 89 ºC, whereas that of potassium is 63 ºC. (CLO3, C4)

HOUR: 10

41. CO2 and BeH2 are triatomic covalent molecules. Describe in detail the formation of the covalent bonds
in these molecules and explain why CO2 obeys the octet rule while BeH2 does not.
PSPM 2013 / 2014 (CLO3, C4)

42. Explain why BeH2 and SF6 disobey octet rule. UPS 2011 / 2012 (CLO3, C3)

43. Based on the skeletal structure of peroxynitrite ion, OONO−, draw all the possible Lewis structures.

Assign formal charge to each atom in the structure. Determine the most stable Lewis structure for

the ion and explain your answer. PSPM 2008 / 2009 (CLO3, C4)

44. a) Formal charge is a useful guide in determining the best or preferred structure. Explain this
b)
statement using [OCN]− ion as example. PSPM 2013 / 2014 (CLO3, C4)

Draw two possible Lewis structures for sulphate ion, SO42− . Determine the most plausible
Lewis structure. UPS 2011 / 2012 (CLO3, C3)

45. Draw the Lewis structures for ClO − and ClO − ions. State the molecular geometry and predict the
3 4

bond angles in each ion. PSPM 2011 / 2012 (CLO3, C4)

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