PRACTICE MODULE CHEMISTRY SK015

PREFACE

This module is a compilation of exercises in Chemistry
Subject for Semester I Session 2022/23. The purpose of
this module is to guide students step by step on how to
answer the questions using specific learning outcomes
as stated in the Curriculum Specifications. Our approach
focuses on understanding and explaining concepts in
Chemistry subject via Mind Map or I-think Map and other
types of exercises. We believe that, understanding the
flow of reaction will make Chemistry an exciting subject
to study and explore. Acknowledgment and appreciation
of real-life applications will lead to realisation that
Chemistry is an ever-expanding and dynamic field of
study. We also appreciate all the efforts made by the
editors:

NOOR ASMAHAN ABDULLAH, FARHANA UMANAN,
ZANARINA SAPIAI, JULIANAWATI AHMED, SITI
AFIZA BASIR, AFRAH ALING, SITI FADHILAH
AYUB, NIK KHADIJAH NIK SALLEH, SITI SARAH
SAIFULLAH, NORAZREEN MUHAMAD, WAN
ROSILAH WAN LLAH, NOOR NAJIHAH
KAMARUDDIN, SHARIFAH FADTHYAH SYED
BAHARUDDIN, SITI FATIMAH MD SOLLHI, AND NUR
ZARIFAH SYAZANA HAMZAH

Hopefully, this effort will help all the students in
learning Chemistry Subject with the right method in
Semester 1 and hoping that they score in the final
exam (PSPM 1). Thank you

ALL THE BEST AND GOOD LUCK

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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
Faraday constant = 4.18 kJ kg−1 K−1
Planck constant = 4.18 J g−1 K−1
Rydberg constant = 4.18 J g−1 C−1

Ideal gas constant NA = 6.021023 mol−1

Density of water at 25C F = 96500 C mol−1
Freezing point of water
Vapour pressure of water at 25C h = 6.625610−34 J s

RH = 1.097  107 m−1
= 2.18  10−18 J

R = 8.314 J mol-1 K−1
= 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
1mL = 1 cm3

ENERGY 1J = 1 kg m2 s−2 = 1 N m= 1  107 erg
1 calorie
1eV = 4.184 J
= 1.602 x 10-19 J

PRESSURE 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

LIST OF CONTENT Page
1-40
No Chapter 41-62
1 Matter 63-74
2 Atomic Structure 75-102
3 Periodic Table 103-130
4 Chemical Bonding 131-156
5 States of Matter 157-186
6 Chemical Equilibrium
7 Ionic Equilibria

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1.1: Atoms and Molecules PRACTICE MODULE CHAPTER 1 SK015

CHAPTER 1 : MATTER

Mind map / I-Think Map

WORKSHEET 1.1
Isotopic Notation

1. Define isotope
_____________________________________________________________________

2. Here are three isotopes of an element:
162 163 164

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PRACTICE MODULE CHAPTER 1 SK015

a. The element is _____________
b. The number 6 refers to the _________________________________
c. The numbers 12, 13, and 14 refer to the ________________________
d. How many protons and neutrons are in the first isotope? _______________
e. How many protons and neutrons are in the second isotope? ______________
f. How many protons and neutrons are in the third isotope? ___________________

3. Give the isotope notation for the isotope of potassium with A=40

4. Write down the isotope symbols of oxygen with proton number of 8 and nucleon number of
16.

5. Complete the table below using suitable answer

Atom Protons Neutrons Electrons Nucleon Charge Isotopic
82 Number Notation
Pb 34 +2
Se 9 207 0 20872Pb2+
Cr 79
F 46 28 21 52 0
Nb 10 4913Nb5+
P 106 93
Rb 80 36 31 -3 3151P3−
Pd 54 85 +1
Os 36 106
K 0
Mo 39
Sg 114 72 96 0
Fr 20 +6
Hg 36 223 0
Xe 159 0
87
121 78
77 54

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PRACTICE MODULE CHAPTER 1 SK015

Interpret Mass Spectrum, Calculate Average Atomic Mass & Relative Atomic Mass

1. Based on the mass spectrum of zirconium above, calculate the average atomic mass of
zirconium. (91.3 amu)

2. Naturally occurring chlorine is 75.78% 35Cl, which has an atomic mass of 34.969 amu, and
24.22% 37Cl, which has an atomic mass of 36.966 amu. Calculate the average atomic mass of
chlorine. (35.45 amu)

3. Naturally occurring magnesium has the following isotopic abundances:

Isotope Abundance (%) Atomic Mass (amu)

24Mg 78.99 23.985
25Mg 10.00 24.985
26Mg 11.01 25.983

a) What is the average atomic mass of Mg? (24.30 amu)

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PRACTICE MODULE CHAPTER 1 SK015

b) Based on data from table above, sketch the mass spectrum of Mg.

4. Boron has two naturally occurring isotopes with the natural abundances shown in the table
below:
Isotope Natural abundance (%)
10B 19.9
11B 80.1
Calculate the relative atomic mass of boron.

5. Lithium has two naturally occurring isotopes: 6Li (7% abundance) and 7Li (93% abundance).
Calculate the relative atomic mass of lithium.

6. Rubidium has a relative atomic mass of 85.47 and consists of two naturally occurring
isotopes, 85Rb (Mass = 84.91 amu) and 87Rb (Mass = 86.91 amu). Calculate the percentage
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PRACTICE MODULE CHAPTER 1 SK015

composition of these isotopes in a naturally occurring sample of rubidium. [Ans : 85Rb = 72%;
87Rb = 28% ]

7. Iridium has a relative atomic mass of 192.22 and consists of Ir-191 and Ir-193 isotopes.
Calculate the percentage composition of a naturally occurring sample of iridium.

[Ans = 191Ir = 39% and 193Ir = 61% ]

1.2 MOLE CONCEPT
I-think map (Definition)

Formula Formula

Simplest Empirical Atoms Exact Molecular atoms of
ratio formula of the numbers formula each
elements
elements

compound molecules

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PRACTICE MODULE CHAPTER 1 SK015

Concentration Measurement (I-Think Map)

Molarity, Percentag
M (mol L- e by

1) mass, w/w
(%)

Molality, Concentratio Percentag
m n e by

(mol Kg-1) Mole volume,
fraction of v/v (%)

A, XA.

*Note: You can write formula/definition in
the blank box above.

WORKSHEET 1.2

Empirical Formula & Molecular Formula

1. Define empirical formula.
___________________________________________________________________________

___________________________________________________________________________
2. Define molecular formula.

___________________________________________________________________________

___________________________________________________________________________

3. What is the empirical formula of a compound that contains 0.783 g of C, 0.196 g of H and 0.521 g of O?
(RAM C = 12, H = 1, O = 16) [Ans : C2H6O]

Substance Carbon Hydrogen Oxygen
Mass (g)
Number of moles (mol)

Smallest ratio

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PRACTICE MODULE CHAPTER 1 SK015

Empirical formula

4. An organic compound that consist of 58.8% carbon, 9.8% hydrogen and 31.4% oxygen has a molar
mass of 102 g mol-1. Determine the empirical and molecular formula of the compound. (Ans :C5H10O2)

Assume mass of compound = 100g

Element C H O

Mass (g) Molecular formula

Number of moles
(mol)

Smallest ratio

Ratio

Empirical formula

5. Compound Y is made up of 39.52% of carbon, 13.14% hydrogen and nitrogen.

i. Determine the empirical formula of compound Y. (Ans : CH4N)

Assume mass of compound = 100g

Element Carbon Hydrogen Nitrogen

Mass (g)

Number of moles

(mol)

Smallest ratio

Empirical formula

ii. the molar mass of compound Y is 30 g mol-1. Determine its Molecular formula. [Ans : CH4N]

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PRACTICE MODULE CHAPTER 1 SK015

6. What is empirical formula if compound consist of 21.2% N, 6.1% H, 24.2% S and 48.5% O? (RAM: N
=14, H= 1, S =32, O= 16). (Ans : N2H8SO4)

Assume mass of compound = 100g

Element Nitrogen Hydrogen Sulphur Oxygen

Mass (g)

No. of moles (mol)

Simplest ratio

Empirical formula

7. Determine the molecular formula of hydrocarbon which contains 85.6% carbon and has a molar mass of
84 g mol-1. (Ans : C6H12)

Assume mass of hydrocarbon = 100 g

Element C H

8. 1.44 g sample of an oxide of copper contains 1.28 g of copper. What is the empirical formula of this
oxide of copper? (Ar Cu= 64, O=16). (Ans : Cu2O)

Element Cu O
Mass (g)

Number of moles (mol)

Smallest ratio
Empirical formula

9. A sample of hydrated salt was found to contain 12.00g Cu, 6.00 g of S, 12.00g of O and 17.00g of
water. Determine the empirical formula. [Ar : Cu = 63.5, O= 16, S=32, H=1] ()

Element Cu S O H2O

Mass (g)

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PRACTICE MODULE CHAPTER 1 SK015

No. of moles (mol)

Simplest ratio

Empirical formula

Empirical and molecular formula involving combustion.

1. After combustion with excess oxygen, a 12.501 g of a petroleum compound produced 38.196 g of
carbon dioxide and 18.752 of water. A previous analysis determined that the compound does not contain
oxygen. Establish the empirical formula of the compound. (Ans : C5H12)

Mass of carbon = __________________________________
Mass of hydrogen = ________________________________

Element Carbon Hydrogen
Mass (g)
Mol (mol)

Simplest ratio

Simplest mol
Empirical formula

2. Combustion of 0.202 g of an organic compound produce 0.361 g CO2 and 0.147 g H2O. Calculate the
empirical formula for this compound. [Ar: C=12.01, H=1, O=16.00]. (Ans : C3H6O2).

Mass of carbon = ___________________________
Mass of hydrogen = _________________________

Element C H O

Mass (g)
Number of moles
(mol)
Smallest ratio

Ratio

Empirical formula

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PRACTICE MODULE CHAPTER 1 SK015

3. A 0.1 g sample of ethyl alcohol known to contain only carbon, hydrogen and oxygen, was burnt
completely in excess oxygen to form the products 0.1910g CO2 and 0.1172 g H2O. What is the empirical
and molecular formula of the compound if the molecular mass is 138 g mol-1. State the number of
hydrogen atoms per molecule present in the above sample. (Empirical formula = C2H6O, Molecular
Formula= C6H18O3, H=18 atoms)

Mass of carbon = ____________________________
Mass of hydrogen = ____________________________
Mass of oxygen = _________________________________

Element C H O

Mass (g)
Number of moles
(mol)
Smallest ratio

Empirical formula

Molecular formula
Number of H atoms

4. Sometimes, athletes illegally use anabolic steroids to increase muscle strength. A forensic chemist
analyzes some tablets suspected of being a popular steroid. He determined that the substance in the tablets
contains only carbon, hydrogen, and oxygen in and has a molar mass of 300.14 g/mole. When a 1.200 g
sample of this study by combustion analysis, 3.516 g of CO2 and 1.007 g of H2O are collected. What is the
molecular formula for the substance? (Ans : C22H28O2)

Mass of carbon = _______________________________
Mass of hydrogen = _____________________________
Mass of oxygen = _______________________________

Substance Carbon Hydrogen oxygen
Mass (g)
Mol (mol)

Simplest mol

Empirical formula
Molecular formula

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PRACTICE MODULE CHAPTER 1 SK015

Concentration Measurement
Molarity

1. Calculate the molarity of a solution made by dissolving 23.4 g of sodium sulphate Na2SO4 in
enough water to form 125 mL of solution? ( Ans :1.32 M).

2. Calculate the molarity of each of the following solutions:
a. 16.4 g CaCl2 in 0.614 L solution. (Ans : 0.24 M)

b. 48.0 mL of 6.00 M H2SO4 diluted to 0.250 L. (Ans : 1.15 M)

3. Calculate the mass of sodium carbonate, Na2CO3 required to prepare 250 mL of 0.5 M solution.
(Ans : 13.25g )

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PRACTICE MODULE CHAPTER 1 SK015

4. What volume (in mL) of sulphuric acid which concentration is 2M and contain 12.6 g of the acid?
(Ans : 64.2 mL)

5. Calculate the mass of sodium hydroxide that needs to be added to 150 g of water to produce 1.2
molal of solution? (Ans ; 7.2g)

6. What is the molal concentration of a solution prepared by dissolving 0.30 mol of CuCl2 in 40.0 mol
of water? (Ans : 0.42mol/kg)

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PRACTICE MODULE CHAPTER 1 SK015

7. Calculate the molality of the following aqueous solutions, 0.840 M sugar (C12H22O11) solution
(density=1.12 g/mL) (1.01 molal)

8. A solution containing 121.8 g of Zn(NO3)2 per litre has a density of 1.107 g/mL. Calculate its molal
concentration. (0.653 mol/kg)

Percentage by mass
1. Calculate the percent by mass of the solute in a aqueous solution containing 5.50 g of NaBr in 78.2
g of solution? [ Ans : 7.03%]

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PRACTICE MODULE CHAPTER 1 SK015

2. Hydrochloric acid can be purchased as a solution of 37% HCl. What mass of this solution contains
7.5 g of HCl? (Ans : 20.27 g)

3. A solution contains 66% H2SO4 by weight and has density of 1.58 g mL-1. How many moles of the
acid present in 1.0 L of the solution? (Ans : 10.64mol)

4. Calculate the amount of water(in gram) that must be added to 5.00 g of urea, (NH2)2CO in the
preparation of 16.2 percent by mass. ( Ans : 25.9g)

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PRACTICE MODULE CHAPTER 1 SK015

5. The density of 20% by mass of nitric acid solution HNO3 is 1.11 g/mL. Calculate the volume acid
needed to prepare 2.0 L of 3.00 mol L-1 HNO3 solution. [1702.7 mL]

6. A mixture containing benzene and toluene has 18.4 g of toluene and its percentage composition is
30%. Calculate the number of moles of benzene in this solution. (Ans : 0.55 mol)

Percentage by volume
1. Calculate the percentage by volume (% v/v) of ethanol in a solution containing 18.6 cm3 of
ethanol in 120 cm3 solution. (Ans :15.5% )

2. Rubbing alcohol is commonly used as an antiseptic for a small cuts. It is sold as 70% by volume
solution of isopropyl alcohol in water. Calculate the volume of isopropyl alcohol use to make
750 mL of rubbing alcohol? (Ans : 525 mL)

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PRACTICE MODULE CHAPTER 1 SK015

Mole fraction
1. What is the mole fraction of CuCl2 in a solution prepared by dissolving 0.30 mol of CuCl2 in 40.0
mol of H2O? (Ans : 0.007)
2. A solution is prepared by mixing 55g of toluene, C7H8 and 55 g of bromobenzene C6H5Br. What
is the mole fraction of each component? (Ans : 0.63, 0.37)

3. Calculate the mole fractions of each compound of the following solutions:
19.4 g of H2SO4 in 0.251 L of H2O (density of water is 1 g/mL) (Ans : H2SO4 =0.014, H2O=0.986)

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PRACTICE MODULE CHAPTER 1 SK015

4. A solution contains 35% by mass HBr and has a density of 1.30 g mL-1. Calculate the molarity
and molality of this solution. (Ans : 5.62 M, 4.32 molal)

5. An aqueous solution of hydrofluoric acid is 30.0% HF, by mass, and has a density of 1.101 g
cm-3. What are the molarity of HF in this solution? [ Ans : 16.5 M]

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PRACTICE MODULE CHAPTER 1 SK015

6. A bottle of commercial sulfuric acid (density 1.787 g/ml) is labelled as 86% by weight. What is
the molarity of acid?[ Ans : 15.7 M]

7. Concentrated phosphoric acid is 90% H3PO4 by mass and the remaining mass is water. The
molarity of H3PO4 in 90% H3PO4 is 12.2 M at room temperature.

i. What density of this solution at room temperature? (Ans : 1.33 g/mL)

ii.What volume (in mL) of this solution is needed to make a 1.00 L of a 1.00 M phosphoric acid?
(Ans : 82.0 mL)

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PRACTICE MODULE CHAPTER 1 SK015

8. A 6.4 molal NaCl solution has a density of 1.2 g cm-3. Calculate
i. the percentage by mass of NaCl. (Ans : 27.24%)

ii. the molarity of the solution [Ar = Na=23, Cl=35.5]. (Ans : 5.588 mol L-1)

9. Calculate the molality of a 6.0 mol/L aqueous solution of an acid HA with a density of 0.878 g
mL-1. [Mr HA = 98.0]. (Ans : 20.69 mol kg-1)

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PRACTICE MODULE CHAPTER 1 SK015

10. Calculate the mol fraction of hydrochloric acid which contain 35% HCl by mass. [Ans : 0.21]

1.3 STOICHIOMETRY Circle Map

LIMITING REACTANT Limiting
1. Limiting reactant is the reactant that reactant
completely used in the chemical reaction
and limit the amount of product formed. By
using definition of limiting reactant,
complete your I-think map(circle map).

Other important formula


= × %

WORKSHEET 1.3

Balancing chemical equation

1. Give oxidation numbers for the underlined atoms in these molecules and ions:

a. NaClO4 _____ b. PtCl62- _____ c. CaI2 _____ d. SnF2 _____

e. Al2O3 _____ f. ClF3 _____ g. H3AsO3 _____ h. SbF6 ̄ _____

i. KNO3 _____ j. MnO2 _____ k. MnO42- _____ l. K2Cr2O7 _____

2. Balance the following chemical equation.
a) Al + O2 → Al2O3
b) Al(NO3)3 + NaOH → Al(OH)3 + NaNO3
c) KNO3 → KNO2 + O2

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PRACTICE MODULE CHAPTER 1 SK015

d) O2 + CS2 → CO2 + SO2
e) KClO3 → KCl + O2
f) BaF2 + K3PO4 → Ba3(PO4)2 + KF
g) H2SO4 + Mg(NO3)2 → MgSO4 + HNO3
h) HCN + CuSO4 → H2SO4 + Cu(CN)2
i) GaF3 + Cs → CsF + Ga
j) N2 + H2 → NH3

3. Balance the following chemical equation. P
a) MgF2 + Li2CO3 → MgCO3 + LiF
b) AgNO3 + Cu → Cu(NO3)2 + Ag
c) AlBr3 + K2SO4 → KBr + Al2(SO4)3
d) C2H6 + O2 → CO2 + H2O
e) NH3 + H2SO4 → (NH4)2SO4
f) K + B2O3 → K2O + B
g) N2 + O2 → N2O5
h) 2NaOH + H2CO3 → Na2CO3 + 2 H2O
i) Al + S8 → Al2S3
j) Ca3(PO4)2 + SiO2 + C → CaSiO3 + 5 CO +

Balancing redox equation

MnO4 - + C2O4 2- → MnO2 + CO2

In Acidic Medium

Step 1: calculate oxidation
number one of the species
to identify which reactant
will undergoes reduction or
oxidation

Step 2: Write incomplete Oxidation:
half equations for the Reduction:
oxidation and reduction
process Oxidation:
Reduction:
Step 3: Balance half Oxidation:
equation in terms of the Reduction:
number of atoms Oxidation:

Step 4: Balance the half
equation in terms of
number of charges by
adding electron.

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PRACTICE MODULE CHAPTER 1 SK015

Step 5: Multiply the half Reduction:
equation with a suitable
factor

*both equation should have same
number of electrons

Step 6: Combine the half
equation to give balanced
overall equation)

In Basic Medium

Step 7: add 8OH- on both
side of chemical equation
to eliminate 8H+. then,
cancel out H2O on both
sides of the equation.

Overall equation in basic
medium

1. Write the balanced half reactions of the following reactions:
a. Cu + NO3– → Cu2+ + N2O4 in acidic solution

b. CO2 + 2 NH2OH CO + N2 + 3 H2O in basic solution

c. 2 H+ + H2O2 + 2 Fe2+ 2 Fe3+ + 2 H2O in acidic solution
d. H+ + 2H2O + 2MnO4- + 5SO2 2 Mn2+ + 5 HSO4- in acidic solution

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PRACTICE MODULE CHAPTER 1 SK015

e. Mn2+ + BiO3- → MnO4- + Bi3+ in acidic solution

2. Use the ion-electron method to complete and balance the following redox equations, occurring in either
acidic or basic aqueous solution, as indicated. Identify the oxidation and reduction half reactions in
each case.
a. In acidic aqueous solution: XeO3 + BrO3– → Xe + BrO4–

b. In acidic aqueous solution: MnO42- + CH3OH → Mn2+ + HCO2H

c. In acidic aqueous solution: Cr2O72– + I- → Cr3+ + IO3–

d. In basic aqueous solution: SO2 + MnO4– → SO 2– + MnO2
4

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PRACTICE MODULE CHAPTER 1 SK015

3. Balance each redox reaction in acid solution.
a. NO3¯ + Fe2+ → HNO2 + Fe3+

b. MnO4 - + S2O3 2- → S4O62- + Mn 2+-

c. ClO3 - + Cl - → Cl2 + ClO2

4. Balance each redox reaction in Basic Solutions
a. Zn + NO3- → Zn(OH)42- + NH3

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PRACTICE MODULE CHAPTER 1 SK015

b. Cu(NH3)42+ + S2O42- → SO32- + Cu + NH3

5. Write balanced equations for the following reactions:
a. Cr(OH)3 + Br2 CrO42- + Br- in basic solution

b. O2 + Sb H2O2 + SbO2- in basic solution

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PRACTICE MODULE CHAPTER 1 SK015

c. ClO2- ClO2 + Cl- in acidic solution

LIMITING REACTANT

1. Suppose 316.0 g aluminum sulfide reacts with 493.0 g of water. Determine limiting reactant. What mass of the
excess reactant remains? Calculate percentage yield of H2S if actual mass of H2S is 197.0 g [Ans : 91.6%]

STEP 1 Al2S3 + 6H2O → 2Al(OH)3 + 3H2S
Determine limiting reactant

Calculate given moles for both reactant

STEP 2 Write relationship between reactants
from the balance chemical equation

STEP 3 Determine mole of H2O needed using
given mole of Al2S3 (or vise versa)

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PRACTICE MODULE CHAPTER 1 SK015

STEP 4 Compare mole H2O needed with mole
H2O given

STEP 5 By using limiting reactant, determine
mole/mass of excess reactant used in the
reaction

STEP 6 Calculate mass of excess reactant
remain

Calculate percentange yield of H2S if actual mass of H2S is 197.0 g
STEP 7 Write formula of percentage yield

STEP 8 Write relationship between limiting
reactant and H2S

STEP 9 Determine mole/mass of H2S based on
mole limiting reactant (theoritical yield)

STEP 10 Determine % yield of H2S

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PRACTICE MODULE CHAPTER 1 SK015

2. Nitric acid react with oxygen gas to give nitrogen dioxide (NO2), a dark brown gas:
2NO + O2 → 2NO2

In one experiment 0.886 mol of NO is mixed with 0.503 mol of O2.
i. Calculate which of the reactant is the limiting reactant.

ii. Calculate the number of mol of NO2 produced.[Ans : 0.886 mol]

3. Consider the reaction:
MnO2 + 4HCl → MnCl2 + Cl2 + 2H2O
If 0.86 mol of MnO2 and 48.2 g HCl react, which reagent will used up first? How many grams of Cl2 will
be produced? ( Ans : 23.4)

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PRACTICE MODULE CHAPTER 1 SK015

4. 2.0 g of aluminum was reacted with solution containing 11.5 g H2SO4 to produce Al2(SO4)3 and
hydrogen gas.
i. Balance the chemical equation.
______________________________________________________________
ii. Determine the limiting reactant

iii. What is the number of mol of hydrogen released?[ Ans: 0.1112 mol]
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PRACTICE MODULE CHAPTER 1 SK015

iv. Calculate then mass of Al2(SO4)3 produced [Ans : 10.85g]

v. Calculate the amount of excess reactant after the reaction completed [Ans:0.59g]

5. Write the balanced equation for the reaction that occurs when iron (II)chloride is mixed with sodium
phosphate forming iron (II) phosphate and sodium chloride.
i. If 23 grams of iron (II) chloride reacts with 41 grams of sodium phosphate, what is the
limiting reagent? How much sodium chloride can be formed? [Ans: 21.21g]

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PRACTICE MODULE CHAPTER 1 SK015

ii. How much of the excess reagent remains when this reaction has gone to completion?
[Ans : 21.18g]

iii. If 16.1 grams of sodium chloride are formed in the reaction, what is the percent yield of this
reaction? [Ans: 75.91%]

6. Zinc and sulphur react to form zinc sulphide according to the equation.
Zn + S ---------> ZnS

If 25.0 g of zinc and 30.0 g of sulphur are mixed,
i. Which chemical is the limiting reactant?

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PRACTICE MODULE CHAPTER 1 SK015

ii. How many grams of ZnS will be formed? [Ans :37.27g]
iii. How many grams of the excess reactant will remain after the reaction is over?[Ans: 17.73 g]
7. Which element is in excess when 3.00 grams of Mg is ignited in 2.20 grams of pure oxygen? What
mass is in excess? What mass of MgO is formed? [Ans : 0.224 g ; 4.98g MgO]

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PRACTICE MODULE CHAPTER 1 SK015

8. When MoO3 and Zn are heated together they react
3Zn(s) + 2MoO3(s) ----------> Mo2O3(s) + 3ZnO(s)

What mass of ZnO is formed when 20.0 grams of MoO3 is reacted with 10.0 grams of Zn? Consider
the reaction. [Ans: 12.45g]

9. Given reaction; I2O5(g) + 5CO(g) → 5CO2(g) + I2(g)
a) 80.0 grams of iodine(V) oxide, I2O5, reacts with 28.0 grams of carbon monoxide, CO.
Determine the mass of iodine I2, which could be produced? [Ans : 50.76g]

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PRACTICE MODULE CHAPTER 1 SK015

b) If, in the above situation, only 0.160 moles, of iodine, I2 was produced.
i) what mass of iodine was produced? [Ans : 40.61g]
ii) what percentage yield of iodine was produced. [Ans : 80%]

10. Silver nitrate, AgNO3, reacts with ferric chloride, FeCl3, to give silver chloride, AgCl, and ferric
nitrate, Fe(NO3)3. In a particular experiment, it was planned to mix a solution containing 25.0 g of
AgNO3 with another solution containing 45.0 grams of FeCl3.
i. Write the chemical equation for the reaction.
_____________________________________________________________________________________________________
ii. Which reactant is the limiting reactant?

iii. What is the maximum number of moles of AgCl that could be obtained from this mixture?
[Ans: 0.1471 mol]

34

PRACTICE MODULE CHAPTER 1 SK015

iv. What is the maximum number of grams of AgCl that could be obtained?[Ans: 21.09 g]

v. How many grams of the reactant in excess will remain after the reaction is over?
[Ans :37.04 g]

11. Solid calcium carbonate, CaCO3, is able to remove sulphur dioxide from waste gases by the
reaction:
CaCO3+ SO2 + other reactants ------> CaSO3 + other products
In a particular experiment, 255 g of CaCO3 was exposed to 135 g of SO2 in the presence of an

excess amount of the other chemicals required for the reaction.
a) What is the theoretical yield of CaSO3? [ Ans : 253.15g]

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PRACTICE MODULE CHAPTER 1 SK015

b) If only 198 g of CaSO3 was isolated from the products, what was the percentage yield of CaSO3
in this experiment?( Ans : 78.21%)

12. A research supervisor told a chemist to make 100 g of chlorobenzene from the reaction of benzene
with chlorine and to expect a yield no higher that 65%. What is the minimum quantity of benzene
that can give 100 g of chlorobenzene if the yield is 65%? The equation for the reaction is:

C6H6 + Cl2 -----------> C6H5Cl + HCl

benzene chlorobenzene

[Ans: 106.67 g]

13. Certain salts of benzoic acid have been used as food additives for decades. The potassium salt of
benzoic acid, potassium benzoate, can be made by the action of potassium permanganate on
toluene.

C7H8 + 2KMnO4 → KC7H5O2 + 2MnO2 + KOH + H2O

toluene potassium benzoate

If the yield of potassium benzoate cannot realistically be expected to be more than 68%, what is the

minimum number of grams of toluene needed to achieve this yield while producing 10.0 g of KC7H5O2?

[Ans : 8.45g]

36

PRACTICE MODULE CHAPTER 1 SK015

WORKSHEET 1.4

MULTIPLE CHOICE QUESTIONS

1. Isotopes differ from each other in what way?
A they have different numbers of neutrons in the nucleus.
B they have different numbers of protons in the nucleus.
C the have different numbers of electrons outside the nucleus.
D more than one response is correct.

2. How many protons, electrons and neutrons are there in an ion of chlorine isotope, 3175 −?
A 35 p, 35 e, 17 n
B 35 p, 17 e, 18 n
C 17 p, 35 e, 17 n
D 17 p, 18 e, 18 n

3. Naturally occurring carbon consists of four isotopes: 11C, 12C, 13C, and 14C. The proton number of
carbon is 6. Which of the following is true about carbon isotopes?
A Attraction forces of 14C nucleus is stronger than of 11C.
B 14C has the most number of neutrons.
C All carbon isotopes are having different atomic size.
D All carbon isotopes contain the same number of neutron and electrons.

4. Which of the following is a pair of isotopes?
A 13588 and 13680
B 7324 and 3735
C3863 and 3846
D 23925 and 23949

5. Which of the following contains only empirical formula?
A C3H8, N2O4, H2O
B C5H12, NO2, H2O
C C5H10, N2O4, H2O2
D C5H12, NO2, H2O2

6. Potassium oxalate has the chemical formula K2C2O4. Based on this information, the formula of
iron(III) oxalate is
A FeC2O4
B Fe(C2O4)2
C Fe2(C2O4)3
D Fe3(C2O4)2

7. An organic compound contains carbon, hydrogen and oxygen. 1.84 g of the organic compound

gives 3.52 g of carbon dioxide and 2.16 g of H2O on complete combustion. Determine the empirical

formula of the organic compound,

A C2H5O C CH4O

B C2H6O D C2H4O

37

PRACTICE MODULE CHAPTER 1 SK015

8. What is the stoichiometric coefficient of H2O when the following equation is properly balanced with
the smallest set of whole numbers?
NH3 + O2 → NO + H2O
A3
B4
C5
D6

9. The redox reaction between Fe3+ and MnO4- is shown as two half equations below:
Fe2+(aq) → Fe3+(aq) + e-

MnO4- (aq) + 8H+(aq) + 5e- → Mn2+(aq) + 4H2O(l)
Calculate the number of moles of MnO4- needed to oxidize 1 mol of Fe2+.

A 0.20 C. 0.50

B 0.40 D 1.00

10. Give S2O82-(aq) + 2e → 2SO42- (aq)
Mn2+(aq) + 4H2O(l) → MnO4- (aq) + 8H+(aq) + 5e
How many moles of S2O82- are needed to oxidise 1 mol of Mn2+?
A 0.4 mol C 2.0 mol

B 1.0 mol D 2.5 mol

11. Calculate the number of hydrogen atoms in 9.60 g of (NH4)2CO3 compound.

A 6.02 x 1022 C 4.82 x 1023

B 3.01 x 1023 D 6.02 x 1023

12. Borax, Na2[B4O5(OH)4], is used as a fireproof insulation material and as a washing powder.
Calculate the mass of boron present in 30 g of borax.
A 4.30 g
B 5.46 g
C 6.00 g
D 6.44 g

13. What volume of water in cm3 should be added to 10.0 cm3 NaOH 6.0 M to produce a solution of

NaOH 0.3 M?

A 10 cm3 C 200 cm3

B 190 cm3 D 500 cm3

14. A student wish to neutralize phosphoric acid with 0.2 M solution of sodium hydroxide. The reaction

equation is as follows:

H3PO4 + 3NaOH → Na3PO4 + 3H2O

What volume of 0.2 M sodium hydroxide is required to neutralize 50 mL phosphoric acid 0.1 M?

A 25 mL C 50 mL

B 75 mL D 150 mL

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PRACTICE MODULE CHAPTER 1 SK015

15. Hydrogen gas can be prepared through the reaction between zinc and hydrochloric acid as follows
Zn(s) + 2HCl (aq) → H2(g) + ZnCl2(aq)
If 13.0 g zinc react with excess HCl, calculate the volume of hydrogen gas in mL produced at
standard temperature and pressure,
A 4.86 C 4450
B 4.45 D 4863

16. Nitrogen monoxide reacts spontaneously with oxygen as represented by the equation: 2NO(g) +
O2(g) → 2NO2(g). in an experiment, 75.0 g of nitrogen monoxide is allowed to react with 64.0 g of
oxygen. Which of the following statement is/are correct?

I Nitrogen oxide is the limiting reactant.

II 115.0g of NO2 is produced.

III 24.0 g of O2 remained unreacted

WORKSHEET 1.5

MODEL OF PSPM QUESTIONS

1. a) Define nucleon number and isotope.

Give the number of protons, neutrons and electrons in each of the following

species.

79

i. 35 Br –

ii. 130 Ba 2+
56

b) A 50 mL of a saturated NaOH solution containing 52% NaOH by weight with a
density of 1.48 g mL-1 is used to prepare a 0.1 M NaOH solution. Determine the
initial concentration of NaOH solution and the amount of water required to prepare

the 0.1 M NaOH solution.

2. a) Compound A consists of the element C, H and O. The complete combustion of
4.624g of the compound A yielded 6.557g of CO2 and 4.026g of H2O.Determine
the empirical formula of compound A.

b) The reaction between acetic acid, CH3COOH and barium hydroxide, Ba(OH)2
produces a salt. Determine the maximum mass of the salt obtained if 17.13g of
barium hydroxide is used.

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PRACTICE MODULE CHAPTER 1 SK015

3. a) FIGURE 1 shows the mass spectrum for copper, Cu, atom which has been
identified to have two isotopes.
FIGURE 1

62.929 64.927
8 8

Relative 2693Cu 65 Cu
abundan 29

62 63 64 65 66
i. Define isotope. Mass/charge

ii. Using the information given in FIGURE 1, calculate the abundance of
each isotope for Cu.

b) Urea, (NH2)2CO is used as fertilizer and in animal feed. It is prepared by reacting

ammonia and carbon dioxide as shown:

2NH3(g) + CO2(g) (NH2)2CO(s) + H2O(l)

In a process, 637.2 g of ammonia is allowed to react with 1142 g of carbon
dioxide.

i. Determine the limiting reagent in the reaction.

ii. Calculate the mass of urea formed.

Prepared by : Mdm Siti Fatimah & Mdm Nur Zarifah Syazana

40

PRACTICE MODULE CHAPTER 2 SK015

CHAPTER 2.0
ATOMIC STRUCTURE

GENERAL OVERVIEW

Bohr's Atomic Describe Bohr's Atomic Model
Model
Explanation existence of energy level in an
Atomic atom
Structure
Energy of elctron

Formation of line spectrum of hydrogen
atom

Formation of Lyman, Balmer, Paschen, Brackett and
Pfund series.

Energy change of an electron during transition

Energy of photon emitted by an electron that prodeuces at
particular wavelength during transition

Perform calculation involving the Ryberg
equation

Ionisation energy of hydrogen atom from
Lyman series

Limitation of Bohr's Atomic Model

de Broglie's postulate and Heisenberg's uncertainty

Quantum Orbital Principle quantum number, n
Mechanics Angular momentum quantum number,l
Four quantum numbers Magnetic quantum number,m
of an electron Electron spin quantum number,s

3-D shapes of s, p and d
orbitals

Electronic Aufbau principle, Hund's Rule and Pauli exclusion
Configuration Principle

spdf notation & orbital diagram

Anomalous elctronic configuration of
Copper and Chromiun

41

PRACTICE MODULE CHAPTER 2 SK015

Worksheet 2.1
1. State four Bohr’s Atomic Model postulate

i) Electron moves in _____________ orbit around the ______________ of an atom.

ii) In the specific __________ level, the energy of electron is __________ in value or
is ___________.

iii) An electron moves in an ___________ energy state will not ________or radiate
energy.

iv) Energy is ___________ or __________ by an electron as it changes from one
allowed energy state to another.

2. Calculate the energy (in J) of an electron when it occupies a level equivalent to the
quantum number of n = 1, n = 2, n = 3 and n = 4.

n=1 n=2

n=3 n=4

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PRACTICE MODULE CHAPTER 2 SK015

3. The series in the hydrogen line spectrum are found in the visible region, infrared region
and ultraviolet region. State the region for each series in the hydrogen line spectrum.

Series in the hydrogen line spectrum Region
Lyman
Balmer
Paschen
Brackett
Pfund

4. i) Show the electron transition (the first four lines) that give rise to Lyman series.

Energy

n=∞
n=5

n=4

n=3

n=2

n=1

Line spectrum : Lyman series

ii) Show the electron transition (the first four lines) that gives rise to Balmer series

Energy

n=∞

n=6

n=5

n=4

n=3

43 n=2

PRACTICE MODULE CHAPTER 2 SK015

Line spectrum :

Balmer series

iii) Show the electron transition (the first four lines) that gives rise to Paschen series.

Energy

n=∞
n=7

n=6

n=5

n=4

n=3

Line spectrum :

Paschen series

5. a) i) A photon is emitted when an electron of hydrogen atom make transition from
energy level, n = 5 to the n = 3. Calculate the energy emitted and the wavelength
corresponding to this transition. [Ans : – 1.55 x 10-19 J] [Ans : 1282 nm]

44

PRACTICE MODULE CHAPTER 2 SK015

ii) A photon is emitted when an electron of hydrogen atom make transition from energy
level, n = 4 to the n = 2. Calculate the energy emitted and the wavelength
corresponding to this transition. [Ans : – 4.0875 x 10-19 J ; 486.3 nm]

b) Calculate the energy emitted and wavelength produced when an electron in
hydrogen atoms is transferred from n = 6 to n = 2.
[Ans : – 4.844 x 10-19 J; 410.3 nm]

6. a) i) Calculate the wavelength and frequency of second line in Lyman series.
[ Ans : 1.026 x 10-7 m @ 102.6 nm ; 2.92 x 1015 s-1]

45