Chemistry Tutorial Book

CHEMISTRY

Tutorial Book

EC015

EC015: CHEMISTRY 1

Distribution of Student Learning Time (SLT)

Teaching and Learning Activities

Course Content Outline Guided Learning Independent TOTAL
SLT
Learnig
LECTURE TUTORIAL PRACTICAL (NF2F)

1.0: MATTER 2 18 6 20 46
1.1 Atoms and molecules
1.2 Mole concept
1.3 Stoichiometry

2.0: ATOMIC STRUCTURE 3 6 0 9 18
2.1 Bohr’s atomic model 2 8 0 10 20
2.2 Quantum mechanics 3 10 0 13 26
2.3 Electronic configuration
3 7 2 10 22
3.0: PERIODIC TABLE 2 7 2 9 20
3.1 Classification of elements
3.2 Periodicity

4.0: CHEMICAL BONDING
4.1 Lewis structure
4.2 Molecular shape and
polarity
4.3 Orbital overlap and
hybridisation
4.4 Intermolecular forces
4.5 Metallic bond

5.0: STATE OF MATTER
5.1 Gas
5.2 Liquids
5.3 Solids
5.4 Phase diagram

6.0: CHEMICAL EQUILIBRIUM
6.1 Dynamic equilibrium
6.2 Equilibrium constants
6.3 Le Chatelier’s Principle

7.0: IONIC EQUILIBRIA 3 10 2 13 28
7.1 Acids and bases 18 66 12 84 180
7.2 Acid-base titrations
7.3 Solubility equilibria

TOTAL

TABLE OF RELATIVE ATOMIC MASSES

Element Symbol Proton number Relative atomic mass
Aluminum Al 13 27.0
Silver Ag 47 107.9
Argon Ar 18 40.0
Arsenic As 33 74.9
Gold Au 79 197.0
Barium Ba 56 137.3
Beryllium Be 4 9.0
Bismuth Bi 83 209.0
Boron B 5 10.8
Bromine Br 35 79.9
Iron Fe 26 55.9
Fluorine F 9 19.0
Phosphorus P 15 31.0
Helium He 2 4.0
Mercury Hg 80 200.6
Hydrogen H 1 1.0
Iodine I 53 126.9
Cadmium Cd 48 112.4
Potassium K 19 39.1
Calcium Ca 20 40.1
Carbon C 6 12.0
Chlorine Cl 17 35.5
Cobalt Co 27 58.9
Krypton Kr 36 83.8
Chromium Cr 24 52.0
Copper Cu 29 63.6
Lithium Li 3 6.9
Magnesium Mg 12 24.3
Manganese Mn 25 54.9
Sodium Na 11 23.0
Neon Ne 10 20.2
Nickel Ni 28 58.7
Nitrogen N 7 14.0
Oxygen O 8 16.0
Platinum Pt 78 195.1
Lead Pb 82 207.2
Protactinium Pa 91 231.0
Radium Ra 88 226.0
Radon Rn 86 222.0
Rubidium Rb 37 85.5
Selenium Se 34 79.0
Cerium Ce 58 140.1
Caesium Cs 55 132.9
Silicon Si 14 28.1
Scandium Sc 21 45.0
Tin Sn 50 118.7
Antimony Sb 51 121.8
Strontium Sr 38 87.6
Sulphur S 16 32.1
Uranium U 92 238.0
Tungsten W 74 183.9
Zinc Zn 30 65.4

LIST OF SELECTED CONSTANT VALUES

Ionization constant for water at 25 C Kw = 1.0 10 14 mol2 dm 16
Molar volume of gases Vm = 22.4 dm3 mol 1 at STP

Speed of light in a vacuum = 24 dm3 mol 1 at room temperature
Specific heat of water c = 3.0 108 m s 1

Avogadro’s number = 4.18 kJ kg 1 K 1
Faraday constant = 4.18 J g 1 K 1
Planck constant = 4.18 J g 1 C 1
Rydberg constant NA = 6.02 1023 mol 1
F = 9.65 104 C mol 1
Ideal gas constant h = 6.6256 10 34 J s
RH = 1.097 107 m 1
Density of water at 25 C = 2.18 10 18 J
Freezing point of water R = 8.314 J mol-1 K 1
Vapour pressure of water at 25 C = 0.08206 L atm mol 1 K 1

= 1 g cm 3

= 0.00 C

P H2O = 23.8 torr

UNIT AND CONVERSION FACTOR

VOLUME 1 L = 1 dm3
ENERGY 1mL = 1 cm3
PRESSURE
1J = 1 kg m2 s 2 = 1 N m= 1 107 erg
1 calorie
1eV = 4.184 J
= 1.602 x 10-19 J

1 atm = 760 mmHg=760 torr =101 325 Pa = 101.325 kPa =101 325 N m-2

OTHERS 1 faraday (F) = 96 500 C
1 newton (N) = 1 kg m s 2

Chapter 1
Matter

CHAPTER 1.0: MATTER

1.1 Atoms & Molecules

1. Write the isotope notation for each of the following species:

Species protons Number of
1 neutrons electrons
A 1
B 9 02
C 2 10
D 10 10
22

2. Determine the number of protons, neutrons and electrons in the following species.
a) 4200

b) 168 2−

c) 2131 +

3. The following is the mass spectrum of zirconium. Calculate the average atomic mass of

zirconium. (91 u)

% abundance intensity
52

9 12 14 13
mass( m/z) a.m.u

90 91 92 93 94

4. Naturally occurring bromine is composed of two isotopes, 79Br with a mass of 78.91834u

and 81Br with a mass of 80.9163 u. Average atomic mass of bromine is 79.90 u. Calculate

the percent abundance of 79Br and 81Br. ( 79 Br = 50.87%, 81Br = 49.13%)

5. Chlorine isotopes occur naturally as 35Cl and 37Cl. The abundance ratio of these two

isotopes is 35 = 3.127. Based on the scale of carbon-12, the relative mass of 35Cl and
37

37Cl are 34.9689 and 36.9659 respectively. Calculate the relative atomic mass of chlorine.

(35.45)

6. Iron consists of 5.82% 54Fe, 91.66% 56Fe, 2.19% 57Fe and 0.33% 58Fe. The isotopic masses

of these four isotopes are 53.9396 u, 55.9394 u, 56.9354 u and 57.9333 u respectively.

Calculate the relative atomic mass of iron. (55.85)

1

CHAPTER 1.0: MATTER

1.2 Mole Concept

7. Calculate the number of moles for: (1 mol H atoms)
a) 6.02×102 3 atoms of hydrogen. (5.0010-3mol O3)
b) 3.01×1021 molecules of ozone. (3.0010-2 mol N2)
c) 1.806×1022 molecules of nitrogen. (0.250 mol H2SO4)
d) 24.5 g of hydrogen sulphate. (0.446 mol O2)
e) 10.0 L O2 gas at STP.

8. Quinine C20H24N2O2, is a compound extracted from cinchona tree which is traditionally

used to treat malaria. If given a 1.08 g of quinine sample, calculate:

a) the molecular mass of quinine. (324.0g mol–1)

b) the number of moles of quinine. (3.3310-3 mol)

c) the number of molecules of quinine. (2.001021 molecules)

d) the number of hydrogen atoms. (4.80 1022 H atoms)

e) the mass of carbon atoms in gram. (0.799 g @ 0.797 g)

9. Analysis of a gaseous hydrocarbon compound gives the following percent composition by

mass: 85.7% C and 14.3% H.

a) Define empirical formula and molecular formula.

b) Determine the empirical formula of the hydrocarbon. (CH2)

c) 0.25 g of this compound occupies a volume of 100 mL at STP. Determine the molar

mass and the molecular formula of the hydrocarbon. (56 g mol–1, C4H8)

10. 0.6216g sample of a compound of C, H, N and O was found. The compound contains

0.1735g C, 0.01455g H, 0.2024g N and have a relative molecular mass of 129. Calculate the

empirical and molecular formula of the compound, arranging their symbols in alphabetical

order. (CHNO, C3H3N3O3)

11. A complete combustion of a hydrocarbon forms 1.10 g of CO2 and 0.45 g of H2O. The
molar mass is 84.00 g mol–1. Determine the empirical formula and molecular formula of

the hydrocarbon. (CH6, C6H12)

12. Saline solution is prepared by dissolving 9g of NaCl in deionised water in a 500 mL

volumetric flask. Calculate the molarity of the solution. (0.3076M)

13. A student would like to prepare 250mL of 0.25molL-1 lead (II) nitrate, Pb(NO3)2 solution.

a) Calculate the mass of Pb(NO3)2 required to prepare the solution (20.7g)

b) Based on the information in (i), give TWO main steps that should be carried out to

prepare the solution.

2

CHAPTER 1.0: MATTER

14. The density of 10.5 molal NaOH solution is 1.33 g mL– 1 at 20°C. Calculate: (0.159)
a) the mole fraction of NaOH. (29.6%)
b) the percentage by mass of NaOH. (9.83 M)
c) the molarity of the solution.

15. The density of 95% by mass of sulphuric acid, H2SO4 is 1.84 g mL-1. Calculate

a) the molarity of H2SO4. (18 M)

b) the volume of the acid needed to prepare 1.0 L of 0.080 M solution. (4.310–3 L @ 4.3 mL)

1.3 Stoichiometry

16. Determine the oxidation number of the underlined elements in the following compounds:
a) NO2
b) KMnO4
c) HClO3
d) H2SO4
e) Cr2O72-
f) IO3-

17. Balance each of the following equations:
a) NaOH(aq) + FeCl3(s) → Fe(OH)3(s) + NaCl(aq)
b) C4H10(g) + O2(g) → CO2(g) + H2O(l)
c) Fe(s) + H2O(l) → Fe3O4(s) + H2(g)
d) Fe2O3(s) + HCl(aq) → FeCl3(aq) + H2O(l)
e) Cr(OH)3(aq) + IO3-(aq) → CrO32-(aq) + I-(aq) (acidic medium)
f) Cl2(aq) → ClO4–(aq) + Cl–(aq) (basic medium)
g) I- + MnO4- + H+ → I2 + Mn2+ + H2O
h) Cl2 + OH- → Cl- + ClO3- + H2O

18. The reaction of powdered aluminium and iron (III) oxide, 2Al + Fe2O3 → Al2O3 + 2Fe

produces so much heat the iron that forms is molten. Because of this, railroads use the

reaction to provide molten steel to weld steel rails together when laying track. Suppose

that in one batch of reactants 4.20 mol of Al was mixed with 1.75 mol of Fe2O3.

a) Which reactant was the limiting reactant?

b) Calculate the number of grams of iron that can be formed from the mixture of

reactants. (196g Fe is formed)

3

CHAPTER 1.0: MATTER

19. A reaction between 7.0 g of copper(II) oxide and 50 mL of 0.20 M nitric acid produces

copper(II) nitrate, Cu(NO3)2 and water.

a) Write the balanced chemical equation for the above reaction.

b) Determine the limiting reactant.

c) Calculate the mass of excess reactant after the reaction. (6.6g)

d) Determine the percentage yield if the actual mass of copper(II) nitrate obtained

from the reaction is 0.85g (90%)

20. A sample of 1.55 g of iron ore is dissolved in an acid solution in which the iron is

converted into Fe2+. The solution formed is then titrated with KMnO4 which oxidises Fe2+
to Fe3+ while the MnO4- ions are reduced to Mn2+ ions. 92.95 mL of 0.02 M KMnO4 is

required for the titration to reach the equivalence point.

a) Write the balanced equation for the titration.

b) Calculate the percentage of iron in the sample. (33.5%)

PSPM 2019/2020

1. (a) Bromine has two stable isotopes, 79Br and 81 Br.

(i) By comparing the number of sub-atomic particles of these isotopes,

explain what is meant by the term isotopes.

(ii) Determine the number of electrons of Br- ion. [4 marks]

(b) In a complete combustion, 1.00g sample W (CxHyOz) was burnt to produce

2.52g of carbon dioxide, CO2 and 0.443g of water vapour, H2O. Determine the

empirical formula of the compound. [6 marks]

(c) Calcium acetate, Ca(C2H3O2)2 solution is the substance used for reducing

phosphate level in late-stage kidney failure. In an experiment, 250mL of

0.25M Ca(C2H3O2)2 solution was prepared. Determine the molality of the

solution with density of 1.509g mL-1. [7 marks]

(d) Magnesium hydroxide, Mg(OH)2 is an antacid that is used to relieve

indigestion, sour stomach and heartburn. It can be prepared by reacting

magnesium chloride, MgCl2 and sodium hydroxide, NaOH with the by-product

of sodium chloride, NaCl. In an experiment, a student allowed 15.1g of MgCl2

to react with 9.35g of NaOH. Calculate the mass (in grams) of Mg(OH)2 that

could be obtained at the end of the experiment. [9 marks]

4

CHAPTER 1.0: MATTER

PSPM 2020/2021 (A)

1. (a) Sodium metasilicate, Na2SiO3, is used in the production of silica gel.

(i) Write the isotopic notation for silicon atom.

(ii) Calculate the number of Si atoms in 50g of Na2SiO3. [3 marks]

(b) The following reaction takes place in an acidic condition.

Write a balance equation for the above reaction. [3 marks]

(c) Compound X has a molar mass of 294.20g. Analysis of X shows that it contains

26.58g of potassium, 35.35g of chromium and 38.07g of oxygen. Determine

the molecular formula and name X. [6 marks]

(d) The reaction between nitrogen dioxide, NO2, with water produces nitric acid,

HNO3, and nitrogen monoxide, NO. In a reaction between 100.0g of NO2 and

50.0g of water, 80.0g of HNO3 is produced.

(i) Determine the limiting reactant in the reaction.

(ii) Determine the percentage yield of the reaction. [9 marks]

PSPM 2020/2021 (B)

1. (a) A species of element X has 54 electrons, 56 protons and 81 neutrons. Write

the isotopic notation for this species. [3 marks]

(b) An aqueous solution of 3.58 m potassium bromide, KBr, with a density of
1.12 g cm−3 is prepared by adding a certain amount of KBr to 100 g of water.
Calculate the molarity and the percentage by mass (% w/w) of the solution.
[8 marks]

(c) An aqueous solution comprising of 0.695 g hydrated copper(II) sulphate,

CuSO4.nH2O, undergoes a complete redox reaction with 0.156 g iron, Fe, to

form iron(II) sulphate, FeSO4 and copper, Cu.

(i) Write the balanced chemical equation and state the changes in

oxidation numbers of Fe and Cu for the redox reaction.

(ii) Calculate the mass of CuSO4 that has completely reacted with the Fe.

(iii) Calculate the hydration value of water, n. [10 marks]

5

Chapter 2
Atomic

Structure

CHAPTER 2.0: ATOMIC STRUCTURE

2.1: Bohr’s Atomic Model

1. (a) Explain why the emission spectrum of hydrogen atom is a line spectrum.
(b) How is the second line of Brackett series produced?

2. The second line of Balmer series has a wavelength of 486.17 nm.

(a) Calculate its frequency. (6.17 x 1014 s-1)
(b) Describe how the line is produced?

3. With reference to the table below:

Colour of line Wavelength /nm
spectrum
Red 656.3
486.3
Green 434.2

Blue

(a) Calculate the energy emitted in the formation of red line and green line.
(3.03 x 10-19 J; 4.09 x 10-19 J)

(b) Sketch the energy level diagram to show the transitions of electron that
produced the red, green and blue lines.

(c) Draw the line spectrum to show the red, green and blue lines.

4. Calculate the wavelength (in nm) and frequency of the fourth line in Lyman series.
(94.96nm; 3.16 x 1015 s-1)

5. In the hydrogen atom, an electron transition from n=5 to a lower energy level
emits photon with a wavelength of 1282 nm.

(a) Determine the lower energy level of this transition
(b) State the series and the radiation region for this transition.

6. When an electron makes a transition from a higher energy level to a lower energy
level, a photon with frequency 6.6 x 1015 s-1 is emitted. Calculate the;

(a) wavelength of the photon. (45.4 nm)

(b) energy emitted by 1 mole of electrons for the above transition.
(2632.48 kJ mol-1)

7. By using the Lyman series, calculate the ionisation energy of hydrogen atom in kJ

mol-1. (1312.4 kJ mol-1)

6

CHAPTER 2.0: ATOMIC STRUCTURE

8. The Lyman series of the spectrum of hydrogen is shown below. Calculate the

ionisation energy of hydrogen from the spectrum. (1312.4 kJ mol-1)

9. Given the spectrum of Lyman series

a b c de
(a) Calculate the energy, wavelength and frequency for line c.

(-2.04 x 10-18 J; 97.2 nm; 3.09 x 1015 Hz)
(b) Which line has the longest wavelength?

2.2: Quantum Mechanics

10. 3d 7s 2d 4p

(a) Which of the above orbitals are allowed?
(b) Determine the maximum number of electrons that can be occupied by

each allowed orbital.
(c) Give sets of quantum number for the allowed orbital.

11. Which of the following quantum numbers are not allowed? Explain your
answer.

(a) (1, 1, 0, +½) (b) (3, 1, -2, +½)
(c) (2, 1, 0, +½) (d) (2, 0, 0, +1)

12. (a) Draw the shape of the following orbitals:
(b) 2s, 3px, 3py, 3dxz, 3dx2-y2, 4dz2
State the principle quantum number and angular momentum quantum
number for each orbital above.

7

CHAPTER 2.0: ATOMIC STRUCTURE

2.3: Electronic Configuration

13. (a) State Aufbau’s principle, Hund’s rule and Pauli exclusion principle.
(b) Arrange the following orbitals in order of increasing
energy.

4dxy, 3dxy, 3dyz, 4pz, 3pz, 3py, 2py, 3s, 2s, 1s, 4s

14. An element Q has proton number of 8. Draw the orbital diagram of the element.
Explain two rules applied in arranging the electrons in the orbital.

15. Write the electronic configuration for the following atoms/ions using spdf notation

andorbital diagram.

i. Cl ii. Cl- iii. Ni iv Ni2+
v. Al vi. Al3+

16. (a) Write the electronic configuration of chromium and copper.
(b) Explain the anomalous electronic configuration in chromium and copper.

17. Given the set of quantum numbers for the highest energy electron in atom Y.
(4, 1, -1, +1/2)

i. Write the electronic configuration for Y.
ii. State the oxidation number of Y ion.
iii. What is the most stable ion of Y? Write its electronic configuration.
iv. Draw the shapes of orbitals occupied by the electrons with the highest

principal quantumnumber in Y

18. The highest energy electrons of element Z occupied the following orbitals in 3rd shell:
zz

yxyx

a. Write the electronic configuration for Z.
b. Write the orbital diagram for its valence electrons.
c. If TWO electrons are being removed from Z, gives the quantum numbers for

these TWO electrons that removed.

8

CHAPTER 2.0: ATOMIC STRUCTURE

PSPM 2019/2020
2. (a) The line spectrum of hydrogen atom in visible region is shown in the
following diagram.

(i) Name the series of the line spectrum.
(ii) Sketch the energy level diagram of a hydrogen atom for the formation

of line X. Explain how line X is formed.
(iii) Calculate the wavelength corresponding to line X.
(iv) Determine the energy involved in 2(a)(iii).

[10 marks]

(b) Describe the anomalous electronic configuration of chromium atom.
[3 marks]

PSPM 2020/2021 (A)

2. (a) Calculate the energy emitted and the wavelength of the spectrum line when

an electron moves from n=6 to n=3. State the series and electromagnetic

region of the spectrum line. [6 marks]

(b) An electron has the following quantum numbers: n=3, l=2. Identify its orbital.

Draw the spatial orientation for all the possible orbitals. [4 marks]

PSPM 2020/2021 (B) [4 marks]
2. (a) Explain why chromium does not obey the Aufbau principle.

(b) Y 2+ ion has 2, 8 and 10 electrons in principal quantum number, n = 1, 2 and 3,

respectively. Determine the atomic number of Y. [2 marks]

(c) An electron in hydrogen atom is excited to n = 5 energy level after it absorbs a

photon of 2.09 × 10 -18 J.

(i) Calculate the energy of this electron.

(ii) Determine the initial energy level of this electron. [4 marks]

9

Chapter 3
Periodic

Table

CHAPTER 3.0: PERIODIC TABLE

3.1: Classification of elements

1. The electronic configuration for atom A-G are as follows,
A: 1s22s22p63s23p4
B: 1s22s22p63s23p63d 84s2
C: 1s22s22p63s23p63d 54s2
D: 1s22s22p63s23p63d 104s24p6
E: 1s22s22p63s23p64s23d104p5
F: 1s22s22p63s23p64s23d104p65s2
G: 1s22s22p63s23p64s23d104p65s24d10

State the i) block, ii) period, iii) group and iv) valence electronic configuration for each
of the atom.

2. Complete the following table: Block electronic configuration
Element Period Group
U 2 17
V 3 15
W46
X 5 17
Y 4 11
Z 62

3.2: Periodicity

3. The following table shows the data for five successive ionization energies of 3 elements,

P, Q and R.

Element Ionization energy(kJ/mol)

IE1 IE2 IE3 IE4 IE5

P 738 1450 7730 10500 13600

Q 1086 2350 4620 6220 38000

R 786 1580 3230 4360 16000

i) Identify two elements which belong to the same group.
ii) Predict the group number of element R
iii) Write the valence electronic configuration for element R

4. The first sixth successive ionisation energy, IE (×103 kJ/mol) of element X is as follows:

IE1 IE2 IE3 IE4 IE5 IE6
1.40 2.86 4.58 7.48 56.27 67.12
Based on the data given, determine the block and group for X. Explain your answer

10

CHAPTER 3.0: PERIODIC TABLE

5. Arrange the following ions in order of increasing ionic radii. Explain your answer.
Mg2+, Al3+, Na+, P3-, S2-, Cl-

6. Sketch the graph of
i. Atomic radius across period 2 from Li to Ne
ii. Ionic radius across period 2 from Li to Ne
iii. 1st Ionisation energy across period 3 from Na to Ar
iv. Atomic radius down the group 1 from Li to Fr
v. 1st ionisation energy down a group 1 from Li to Fr
vi. Successive ionisation energy for 13Al

vii. Electronegativity across period 3 from Na to Ar
viii. Electronegativity down the group 17
Explain the trend of all the graphs
7. The graph shows all the successive ionisation energies of an element P.

Ionisation energy

X
X
XX
X
XX
XX
X
X
X
XX
X XX

No. of electrons removed
i. Determine the group and write the valence electronic configuration for

element P.
ii. Classify the oxide of P and write its equations. State the type of bonding

involved in the formation of oxide P.

11

Chapter 4
Chemical
Bonding

CHAPTER 4.0: CHEMICAL BONDING

4.1 Lewis Structure

1 For each of the following elements, write the electronic configuration for its most
stable cation/anion and state the type of stability:
i. Al : 1s2 2s2 2p6 3s2 3p1
ii. Fe : 1s2 2s2 2p6 3s2 3p6 4s2 3d6
iii. Zn : 1s2 2s2 2p6 3s2 3p6 4s2 3d10
iv. F : 1s2 2s2 2p5
v. N : 1s2 2s2 2p3
vi. S : 1s2 2s2 2p6 3s2 3p4

2 Describe the formation of bonding for each of the following molecules by using Lewis
dot symbol
i. MgBr2
ii. Al2O3
iii. BaO
iv. NH3
v. CO2
vi. N2

3 Describe the formation of covalent dative bond by using these examples:
i. NH4+
ii. Al2Cl6
iii. NH3BF3

4 What is meant by resonance structure? Draw the resonance structures for:
i. NO3-
ii. SO3
iii. N3-

5 (a) Write all possible structures for NCO- ion. Determine the formal
charge of each atom in the structures.

(b) Which structure is the most plausible structure? Explain.

12

CHAPTER 4.0: CHEMICAL BONDING

4.2 Molecular Shape and Polarity
4.3 Orbital Overlapping and Hybridisation

1 For the following molecules
PF3 SO2 HCN CO2
ClF3 SeF4 BrF5 CHCl3

i. Draw the Lewis structure
ii. Predict the molecular geometry according to VSPER theory
iii. Deduce the polarity
iv. Illustrate the hybridisation

2 Illustrate the formation of sigma bond/pi bond from overlapping orbital of the
following compound
i. HBr
ii. N2

3 Illustrate the hybridization of the following molecules
i. NO3-
ii. I3-
iii. NO2+

4 Illustrate the hybridization of the center atom in NH2CH2COOH

4.4 Intermolecular Forces

1 Fill up the table
Structure Polarity Type of Intermolecular Forces
I2
SO2
CH4
H2S
C6H6
NH3

2 (a) Give the factors that influence the strength of Van der Waals forces.
(b) Which molecule has a higher boiling point? Explain.
i. Br2 or ICl
ii. I2 or Br2
iii. ethanol, C2H5OH or dimethyl ether, CH3OCH3

13

CHAPTER 4.0: CHEMICAL BONDING

4.5 Metallic bond

1 (a) Describe the metallic bond in aluminium by using electron sea model.
(b) Explain the electrical conductivity of aluminium.
(c) Explain the why the boiling point of calcium is higher than potassium.

2 Compare the electrical conductivity of sodium and magnesium by using band theory.

PSPM 2018/2019

3. (a) Ammonia NH3 and boron trifluoride, BF3 are covalent compounds. NH3 and BF3 reacts

to form H3NBF3 molecule.

(i) Explain why NH3 obeys the octet rule but BF3 does not.

(ii) Show the formation of H3NBF3 molecule using Lewis dot symbol and label the

bond formed. [5marks]

(b) Oxygen difluoride, OF2 is a strongly oxidizing colourless gas. [8marks]
(i) Determine the molecular geometry of this molecule.
(ii) Explain whether OF2 is a polar or non-polar molecule.

(c) Aluminium and sodium are metals.

(i) Explain the formation of metallic bond in sodium using the electron sea

model.

(ii) Why aluminium has higher boiling point than sodium. [4marks]

PSPM 2019/2020

3. (a) Explain each of the following statements.
(i) Upon reaction with fluorine, oxygen forms only OF2 whereas sulphur forms

SF2. SF4 and SF6 molecules.
(ii) The shape of PF5 molecule differs from that of an IF5 molecule
(iii) Of the three possible resonance structures for OCN- below, III is the best

structure.

[O=C=N]- [O≡C-N]- [O-C≡N]-

I II III

[10 marks]

(b) Illustrate the hybridization of the central atom in SF4 using orbital diagrams.

Show and label the overlapping of orbitals in the molecule. [7marks]

14

CHAPTER 4.0: CHEMICAL BONDING

(c) Explain the difference in melting point between elements in Group1 and Group 17.
[5 marks]

PSPM 2020/2021 (A)

3. Phosgene, COCl2, is a chemical used in the production of plastics and pesticides.
Given that chlorine cannot be the central atom,

(a) draw three (3) possible structures of phosgene.

(b) determine the most plausible structure and give your reason

(c) determine the hybrid orbital for the central atom of the most plausible
structure

(d) name the molecular shape and raw the overlapping orbitals of COCl2
[17 marks]

PSPM 2020/2021 (B)

3. (a) Draw the Lewis structure and name the molecular shape of ClF3 and CF4.

Choose the more polar molecule. [5 marks]

(b) The carboxylate ion, CH3CO2-, has an arrangement of atoms as shown below:

(i) Complete the structure of the carboxylate ion by adding bonds and

lone pair electrons.

(ii) Identify the types of hybridization of C atoms in the ion.

(i) Show and label the possible intermolecular interaction(s) between

CH3COOH molecules. [6 marks]

(c) Calcium is an element in Group 2 of the Periodic Table.

(i) Sketch the electron sea model diagram. [6 marks]
(ii) Describe the type of bonding formed.
(iii) Explain its conductivity.

15

Chapter 5
States of
Matter

CHAPTER 5.0: STATE OF MATTER
5.1: GAS
1. Complete the following table

Gaseous Law Definition Graph Relationship
Boyle’s Law
Charles’ Law
Avogadro’s Law

2. A sample of gas occupies a volume of 10.0 L at the pressure of 2.0 atm. What would be the
pressure of the gas if it is allowed to expand in a 50.0 L container at the same temperature?
(0.4 atm)

3. A sample of gas occupies 100.0 mL at 25 ⁰C. What volume would the gas occupy at 32 ⁰C if
the pressure remains constant?
(102.35 mL)

4. The density of gas X is 2.60 g L–1 at 25 ⁰C and 101 kPa. What is the molecular mass of gas X?
(63.82 g/mol)

5. When an evacuated glass vessel weighing 134.74 g was filled with an unknown gas Y, the

pressure was found to be 99.3 kN m–2 at 31⁰ C, while the mass was 137.28 g. The glass

vessel was then filled completely with water and the mass was 1067.90 g. By using the ideal

gas equation, determine the relative molecular mass of gas Y. (69.32)

6. Figure 1 shows two connecting vessels containing different gases.

When the valve between the vessels is opened, the gases are allowed to mix. Ignoring the

volume taken by the valve, calculate the partial pressure of each gas and the final pressure

of the mixture at 25°C. (3.08 atm, 1.53atm, 4.61 atm)

1L 3L
0.50 mol 0.25 mol

H2 CO2

FIGURE 1

7. Silver oxide, Ag2O when heated, it decomposes to form silver and oxygen gas. Calculate the
volume of gas released at temperature 28.3⁰C and pressure 749 torr if 4.262 g of Ag2O
decomposed completely.
(0.26 L)

16

CHAPTER 5.0: STATE OF MATTER

8.

a) A gas mixture containing 2.45 g of N2 and 3.10 g of Ne occupies a volume of 2.5 L. What

is the pressure of the gas mixture at 25oC. (2.36 atm)

b) 4.0 L of nitrogen at a pressure of 400 kNm-2 and 1.0 L of argon at a pressure of

200 kN m-2 are transferred into a 2.0 L container. Calculate:

i. the partial pressure of nitrogen. (800 kNm-2)

ii. the partial pressure of argon. (100 kNm-2)

iii. the total pressure of the mixture. (900 kNm-2)

9. By using water displacement method, 128 mL of oxygen gas was collected from the

decomposition of potassium chlorate at 24oC and atmospheric pressure of 762 mm Hg.

Calculate the mass of the oxygen gas obtained. The vapour pressure of water at 24oC was

24 mm Hg. (0.163 g)

10. An experiment was carried out to determine the relative molecular mass of CO2 gas. The
following data were obtained.

Mass of empty bulb 25.40 g
Mass of bulb filled with CO2 gas at 17⁰C and pressure of 1.1 × 103 atm 26.50 g
619.50 g
Mass of bulb filled with water

Calculate the relative molecular mass of CO2 gas. (0.04)

There is a difference in value of the relative molecular mass between the one

calculated above and the one calculated from the molecular formula. Explain.

11. The van der Waals equation for n moles of a real gas is

(P + 22 )( − ) =

where a and b are van der Waals constants. The values of a and b for gas A, B and C are
given in the table below

Gas Molar volume/L van der Waals constant/arbitrary unit

ab

A 22.398 0.034 0.02370

B 22.413 1.390 0.03913

C 22.414 3.592 0.04276

a) Arrange the three gases in the above table according to increasing ideality of gases
b) Gas C is more ideal at elevated temperatures. Explain.
c) Almost all gases behave nearly ideal at low pressure. Explain.

17

CHAPTER 5.0: STATE OF MATTER

5.2 : LIQUID

12. Based on kinetic-molecular theory, explain vaporisation and condensation process.

13. Table below shows the vapour pressure of compound A and compound B at room
temperature.

Compound Vapour pressure
(torr)
A 55.3
B 92.0

Which compound has a higher boiling point? Explain.

14. The following table shows the boiling points for a few liquids

Liquid Ethanal (CH3CHO) Ethanol Methanol (CH3OH)
(CH3CH2OH) 65
Boiling Point (⁰C) 20
78

Arrange the liquids in order of increasing strength of intermolecular forces. Explain.

5.3 : SOLID

15. Explain the following processes in terms of phase change and molecular motion.
(a) freezing
(b) melting
(c) sublimation
(d) deposition

16. Using copper, diamond, sulphur and sodium chloride to represent various crystalline
solids, identify the types of interparticle bonding that forms the state of the matter.

17. Give a brief comparison between a solid and a liquid according to the properties
given in Table by filling in the blank.

Properties Solid Liquid
Shape

Surface tension
Viscosity

Compressibility

18

CHAPTER 5.0: STATE OF MATTER
5.4 : PHASE DIAGRAM

18. Figure below shows the phase diagram of water.

a) Which curve represents the equilibrium between ice and water vapour?
b) State the phase changes when a sample at point E is heated at constant pressure

until point F is reached.
c) Name the point at which the BC line intersects 1 atm line.
d) The BC line has a negative slope. Explain.
19. Based on the following data;
Sketch and label the phase diagram. Hence, determine whether liquid A or solid A is
denser. Explain your answer.

Melting point of pure A : 20 ⁰C
Boiling point of pure A : 98 ⁰C
Critical point : 140 ⁰C and 1200 torr
Triple point : 38 ⁰C and 500 torr
20. Carbon dioxide has a triple point at temperature -57 ⁰C and pressure 5 atm while its
critical point is at temperature 31 ⁰C and pressure 73 atm.
Sketch a labeled phase diagram for carbon dioxide. Explain how the melting point of
carbon dioxide changes with increasing pressure.

19

CHAPTER 5.0: STATE OF MATTER

PSPM 2018/2019

4. a) A 10-L cylinder contains 4 g of hydrogen gas and 28 g of nitrogen gas. If the
temperature is 31oC
i. Determine the total pressure of the gaseous mixture
ii. Calculate the partial pressure of hydrogen gas
iii. What will happen if the gaseous mixture is heated to 550oC?
[7 marks]

b) Under the same condition of temperature and density, determine which gas

behaves less ideally: CH4 or SO2 [3 marks]

c) In an experiment when gelatine was added to water, the water became viscous.
Explain the relationship between viscosity and intermolecular forces. [2 marks]

PSPM 2019/2020

4. a) When coal is burnt, the sulphur present in coal is converted to sulphur dioxide
which is responsible for the acid rain phenomenon

S (s) + O2 (g) → SO2 (g)

If 2.54 kg of sulphur is reacted with oxygen, determine the volume of sulphur

dioxide gas formed at 30.5oC and 851.2 mmHg [5 marks]

b) In an experiment, NH3 gas was produced and collected using water displacement
method. At 24oC and atmospheric pressure of 762 mmHg, the volume of the gas

collected was 128 mL. Calculate the mass of the gas obtained.

[Given the vapour pressure of water = 22.4 mmHg] [6 marks]

PSPM 2020/2021

4. a) Sketch a labelled phase diagram of carbon dioxide [4 marks]

b) 2.0 g of He and 61 g of O2 were placed in a 5.0 L tank at 25oC. Determine the partial

pressure and the total pressure of the gas mixture [5 marks]

20

Chapter 6
Chemical
Equilibrium

CHAPTER 6.0 : CHEMICAL EQUILIBRIUM

6.1 Dynamic Equilibrium

1 Nitrogen dioxide (NO2) dimerised reversibly to form dinitrogen tetroxide (N2O4).
i. Write the equilibrium chemical equation for this reaction.
ii. Sketch the curve of concentration against time for this reaction. Explain.

6.2 Equilibrium Constant

2 Write the equilibrium law for the below reactions in terms ofconcentration, Kc,
and/or partial pressure, Kp.
a. HF (aq) ⇌ H+ (aq) + F- (aq)
b. Ag+ (aq) + Fe2+ (aq) ⇌ Ag (s) + Fe3+ (aq)
c. 3NO (g) ⇌ NO2 (g) + N2O (g)
d. Fe3O4 (s) + 4H2 (g) ⇌ 3Fe (s) + 4H2O (l)
e. FeO (s) + CO(g) ⇌ Fe(s) + CO2(g)

3 The following reaction achieved equilibrium when the partial pressure of bromine

gas is 0.60 atm. Determine Kp and Kc. (Kp = 0.775; Kc = 0.157)

FeBr3(s) ⇌ FeBr2(s) + ½ Br2(g)

4 The equilibrium constant Kp for the reaction
2NO2(g) ⇌ 2NO(g) + O2(g)

is 158 at 1000 K. At equilibrium, partial pressures of NO2 and NO are 0.400 atm and
0.270 atm respectively. Calculate the equilibrium partial pressure of O2.

(346.8 atm)

5 At 668 K, 1.00 mol each of CO and Cl2 are introduced into an evacuated 1.75 L flask.
At equilibrium, the total pressure of the gaseous mixture is 32.4 atm. Calculate Kp.
CO(g) + Cl2(g) ⇌ COCl2(g)
(25.0@ 24.96)

6 Bromine gas is allowed to reach equilibrium according to the equation,

Br2(g) ⇌ 2Br(g) Kc = 1.1×10-3 at 1280oC

If the initial concentrations of Br2 and Br are 0.063 M and 0.012 M respectively,
calculate the concentrations of these species at equilibrium.

([Br] = 0.0084 M, [Br2] = 0.00648 M)

7 Ammonium hydrogen sulfide decomposes according to the reaction.

NH4HS(s) ⇌ NH3(g) + H2S(g) Kp = 0.11 at 250 oC .

If 55.0 g of solid NH4HS is placed in a sealed 5.0 L container, what is the partial pressure
of NH3 and H2S at equilibrium?
(0.33 atm)

21

CHAPTER 6.0 : CHEMICAL EQUILIBRIUM

8 If 0.024 mol of N2O4 is allowed to reach equilibrium with NO2 in a 0.372 L flask at 25oC,

N2O4(g) ⇌ 2NO2(g) Kc = 4.61×10-3

calculate the degree of dissociation of N2O4. (0.125)

9 At 430oC, the equilibrium constant, Kc, for the reversible reaction is 4.18×10-3.

H2(g) + I2(g) ⇌ 2HI(g)
If 0.040 mol of HI, 0.01 mol of H2 and 0.030 mol of I2 are initially placed in a 2.0 L
container, is the system at equilibrium? Explain.

10 At 100oC, Kp = 60.6 for the reaction
2NOBr(g) ⇌ 2NO(g) + Br2(g)

In a given temperature, 0.10 atm of each component is placed in a container. Is the
reaction at equilibrium? If not, in which direction will it proceed? Determine whether
more NOBr or Br2 is formed.

6.3 Le Chatelier’s Principle
11 Consider the following equilibrium system:

4NH3(g) + 3O2(g) ⇌ 2N2(g) + 6H2O(g) H = -1531 kJ
Determine in which direction will the reaction shift to re-establish equilibrium if the
system is disturbed by:

i. adding O2.
ii. removing NH3.
iii. increasing the temperature.

12 Decomposition of ethane produces ethene and hydrogen gas.

C2H6(g) ⇌ C2H4(g) + H2(g) H = + ve

By using Le Chatelier’s principle, explain the shift in equilibrium position, if any, of the
reaction if

(a) the concentration of hydrogen gas is decreased.

(b) the temperature is lowered.

(c) a catalyst is added.

(d) C2H6 is removed from the system.

(e) the volume of the container is increased.

(f) the pressure is increased.

(g) an inert gas is added at constant pressure.

(h) an inert gas is added at constant volume.

22

CHAPTER 6.0 : CHEMICAL EQUILIBRIUM

PSPM 2019/2020

5 (a) Sulphur trioxide, SO3 gas decomposes to sulphur dioxide and oxygen gas

according to the following equation:

2SO3(g) ⇌ 2SO2(g) + O2(g)

When 1.0 mol of SO3 is placed into a 2L vessel and heated to 344 K, the system

achieves equilibrium and 0.6 mol of SO3 gas is remained.

(i) Calculate the concentrations of each gas at equilibrium.

(ii) Calculate the equilibrium constant Kc at 344 K. [7 marks]

(b) Phosphorus pentachloride, PCl5 is left in a sealed container to establish
equilibrium.

PCl5(g) ⇌ PCl3(g) + Cl2(g) H = -ve

(i) Explain the effect of lowering temperature on the equilibrium constant,

Kp of the system.

(ii) Explain the effect of adding argon gas at constant volume on the

equilibrium position. [4 marks]

PSPM 2020/2021 (A)

5 (a) Ammonia, NH3, is produced at 500 °C and 200 atm in the presence of Fe2O3

catalyst according to the following equation:

3H2(g) + N2(g) ⇌ 2NH3(g) H = -92.22 kJ

Suggest two (2) changes that can be made to increase the yield of NH3. Explain

your suggestions. [4 marks]

(b) Consider the following reaction:

2SO2(g) + O2(g) ⇌ 2SO3(g) Kc = 4.62

Determine the direction of the reaction if the concentrations of SO2(g), SO3(g)

and O2(g) are 0.2 M, 0.15 M and 8.55 x 10-3 M, respectively. [5 marks]

PSPM 2020/2021 (B)

5 (a) Steam react with powdered charcoal to produce a combustible gas at 400 °C.

H2O(g) + C(s) ⇌ H2(g) + CO(g) H = +206 kJ mol-1

When 1.5 atm of steam was initially used, the pressure of the vessel increased
to 2.4 atm at equilibrium.

(i) Calculate the partial pressure of all gases in the equilibrium mixture.

(ii) Calculate the equilibrium constant, Kc, at 400 °C. [7 marks]

(b) Explain how the equilibrium constant, Kc, will be affected when the volume of

the vessel is reduced to 50%. [2 marks]

23

Chapter 7
Ionic

Equilibria

CHAPTER 7.0 : IONIC EQUILIBRIA

7.1 Acids and Bases
1. Identify the conjugate acid-base pairs in each of the following reactions:

(a) CH3COO− + HCN CH3COOH + CN−

(b) HCO3− + HCO3− H2CO3 + CO32−

(c) H2PO4− + NH3 HPO42− + NH4+

(d) HClO + CH3NH2 CH3NH3+ + ClO−

2. The concentration of OH- in a certain household ammonia cleaning solution is 0.0025 M.

Calculate the concentration of hydronium ion, H3O+. (4.0 x 10-12 M)

3. Calculate the mass of NaOH needed to prepare 500.0 mL of solution with a pH of 10.00?
(0.002 g)

4. The table below shows the base ionisation constant, Kb, for several selected compounds.

Compound Kb
C6H5NH2
N2H4 3.8 x10 -10
NH3 1.7 x10 -6
NH2OH 1.8 x 10 -5
1.1 x 10 -8

a) Arrange the compounds in order of increasing base strength.

b) Give the conjugate acid for each compound and arrange them in order of

increasing acid strength.

5. The table below lists the value of acid dissociation constant, pKa, for some acids at 298 K.

Acid pKa
CH3COOH 4.74
HCN 9.31
HNO2 3.30

a) Calculate the Ka value for CH3COOH and HCN. (1.8x10-5, 4.9x10-10)

b) Arrange the acids in descending order of acid strength.

6. Calculate the pH and the percentage ionization of a 0.35 M solution of propanoic acid,
CH3CH2COOH at 25oC.

[Given Ka CH3CH2COOH = 1.35 x 10- 5] (2.7, 0.57)

7. a) The percentage ionisation of 0.10 M diethylamine solution, (C2H5)2NH. was 9.33 %.
Calculate its Kb.

b) Calculate the pH of a 0.01 M solution of diethylamine, (C2H5)2NH.
(9.6 x 10-4; 11.4)

24

CHAPTER 7.0 : IONIC EQUILIBRIA

8. A sample of 0.050 g of hydrogen fluoride, HF was dissolved in 25.00 mL of water and
titrated with 0.1 M NaOH. The acid required 25.00 mL of the base to reach the
equivalent point. Explain qualitatively the pH of the solution at the equivalent point.

9. Write the hydrolysis equation for the following salts and classify them as acidic, basic or
neutral.

a) NaCN
b) N2H5Cl
c) NaNO3

10. A student is required to prepare a buffer solution containing 0.20 M CH3COOH and

0.30M CH3COONa.

a) Calculate the pH of the buffer solution.
b) Calculate pH change if

i) 0.01 mol of HCl
ii) 0.02 mol of NaOH.
was added to 1.0 L of the buffer solution.

Explain.

[Ka CH3COOH = 1.8x10-5]

(4.92; 0.04; 0.07)

11. A buffer solution is prepared by mixing 400 mL of 1.50 M NH4Cl solution with 600 mL

of 0.10 M NH3.

a) Calculate the pH of the buffer solution. (8.3)

b) Calculate the pH of the buffer solution after the addition of

i. 1 mL of 0.10 M NaOH (8.3)

ii. 1 mL of 0.10 M HCl (8.3)

[Assume that the volume of the solution does not change when HCl and NaOH are

added, Kb NH3 = 1.8 x 10-5]

25

CHAPTER 7.0 : IONIC EQUILIBRIA

7.2: Acid-base Titrations

12. In an experiment, 30 mL of 0.1 M HCl is titrated with 0.1 M NaOH.

a) Calculate the pH of HCl before titration.
b) Calculate the pH of solution after adding 15 mL NaOH.
c) Find the volume of NaOH used and determine the pH at equivalence point.
d) Calculate the pH of solution after adding 50 mL of NaOH.
e) Sketch the titration curve and suggest the indicator.

(1.0; 1.48; 7; 12.4)

13. Sketch and interpret the variation of pH against titrant volume for titrations
between:
a) 40.0 mL of 0.200M NaOH is used to titrate with 0.400 M CH3COOH.
b) 40 mL of 0.10 M HCl is used to titrate with 0.20 M of NH3.

[Ka CH3COOH = 1.8 x10 -5] ; [Kb NH3= 1.8 x10 -5]

Hence, suggest the suitable indicator for the titration above.

7.3: Solubility Equilibria

14. The solubility of silver sulphate is 1.510-2 mol L−1. Calculate the solubility product of the
salt. (1.35 x 10-5)

15. It was found experimentally that the Ksp of calcium sulphate is 2.410−4. Calculate the

molar solubility of calcium sulphate. (1.55 x 10-2)

16. Will precipitate form if 200.0 mL of 0.004 M BaCl2 is added to 600.0 mL of 0.008 M K2SO4?

[Ksp BaSO4= 1.110−10]

17. Calculate the solubility of silver chromate, Ag2CrO4 at 25C in

(a) pure water (1.30 x 10-4)
(b) solution of 0.005 M K2CrO4 solution. (2.12 x 10-5)
[Given Ksp Ag2CrO4 = 9.010−12]
Explain the difference in the values of solubility.

26

CHAPTER 7.0 : IONIC EQUILIBRIA

PSPM 2019/2020
6.

a) Nitrous acid, HNO2 is used to distinguish between primary, secondary and tertiary
amines. At 25 0C, an aqueous solution of 0.215M HNO2 is 4.0% dissociated.

(i) Write a balanced equation for the discociation of HNO2 in water.
(ii) Determine the acid dissociation constant, Ka for HNO2.
(iii) Calculate the pH of the solution.

[6 marks]

(b) Calcium hydroxide, Ca(OH)2 is used to neutralized excess acidity in lakes and soils.
Calculate the pH of 0.30 M Ca(OH)2 solution.
[4 marks]

(c) In a weak acid-strong base titration, 25.00 mL KOH solution, is titrated with 21.45 mL
of 0.10 M HCN to reach the equivalent point.
(i) Write a balanced chemical reaction equation for the titration.
(ii) Calculate the concentration of KOH solution at the equivalent point.
(iii) Explain why potassium cyanide, KCN is not a neutral salt using appropriate
equations.
(iv) Suggest a suitable indicator for the titration.
[7 marks]

PSPM 2020/2021 (A)
6.

(a) Determine the pH of 0.5 M ethanoic acid, CH3COOH solution.
[Given Ka CH3COOH = 1.8 x 10-5]
[5 marks]

(b) In an experiment, 25 mL of CH3COOH was titrated against 0.5 M NaOH solution using
phenolphthalein as indicator. The end point is reached when 25 mL of NaOH solution
was added.
(i) Determine the concentration of the product.
(ii) Determine the pH of the solution at end point
[9 marks]

PSPM 2020/2021 (B)
6.

(a) Boric acid, H3BO3 is a food addictive that effectively inhibit the growth of yeasts. The
pKa value of boric acid is 9.24.
(i) Write the dissociation equation of H3BO3 in water.
(ii) Calculate the pH of a 0.20 M solution of H3BO3.
[9 marks]

(b) Lead sulphate, PbSO4, is a slightly soluble salt with a solubility product, Ksp of
1.6 x 10-8 at 25 0C.
(i) Write the Ksp expression for PbSO4.
(ii) Calculate the solubility, s, of PbSO4 in 0.1 M H2SO4 solution. [5 marks]

27

ACKNOWLEDGEMENTS

A very special thanks and appreciation to all
Chemistry Unit members:

Au Siew Wai
Dr. Usha Devi A/P Ramasundram

Haryati binti Hj Harun
Mohd Zahir Hafizi bin Safian

Norhanida binti Elias
Nurain binti Jamhari
Nur Hamizah binti Ab Rahim
Nurulfarhana binti Khamis
Nurbalqis binti Zailan

Chemistry Unit
Kolej Matrikulasi Kejuruteraan Johor