i
TABLE OF CONTENTS
SYLLABUS OUTLINE ..................................................................................... 1
TABLE OF RELATIVE ATOMIC MASSES..................................................... 2
LIST OF SELECTED CONSTANT values ........................................................ 3
TUTORIAL 1.1 : ATOMS AND MOLECULES ................................................ 4
TUTORIAL 1.2 : MOLE CONCEPT.................................................................. 6
TUTORIAL 1.3 : STOICHIOMETRY................................................................ 9
TUTORIAL 2.1 : BOHR’S ATOMIC MODEL ................................................ 10
TUTORIAL 2.2 : QUANTUM MECHANICAL MODEL ................................ 12
TUTORIAL 2.3 : ELECTRONIC CONFIGURATION .................................... 13
TUTORIAL 3.1 : CLASSIFICATION OF ELEMENTS................................... 14
TUTORIAL 3.2 : PERIODICITY..................................................................... 15
TUTORIAL 4.1 : LEWIS STRUCTURE.......................................................... 18
TUTORIAL 4.2 : MOLECULAR SHAPE AND POLARITY........................... 20
TUTORIAL 4.3 : ORBITAL OVERLAP AND HYBRIDIZATION ................. 21
TUTORIAL 4.4 : INTERMOLECULAR FORCES .......................................... 22
TUTORIAL 4.5 : METALLIC BOND.............................................................. 23
TUTORIAL 5.1 : GAS ..................................................................................... 24
TUTORIAL 5.2 : LIQUID................................................................................ 26
TUTORIAL 5.3 : SOLID.................................................................................. 27
TUTORIAL 5.4 : PHASE DIAGRAM ............................................................. 28
TUTORIAL 6.1 : DYNAMIC EQUILIBRIUM ................................................ 30
TUTORIAL 6.2 : EQUILIBRIUM CONSTANTS ............................................ 31
TUTORIAL 6.3 : LE CHATELIER’s PRINCIPLE........................................... 33
TUTORIAL 7.1 : ACIDS AND BASES ........................................................... 34
TUTORIAL 7.2 : ACID-BASE TITRATIONS................................................. 36
TUTORIAL 7.3 : SOLUBILITY EQUILIBRIA ............................................... 37
HANDOUT 1 ................................................................................................... 38
HANDOUT 2 ................................................................................................... 39
HANDOUT 3 ................................................................................................... 40
HANDOUT 4 ................................................................................................... 41
HANDOUT 5 ................................................................................................... 44
HANDOUT 6 ................................................................................................... 45
SYLLABUS OUTLINE CHEMISTRY (SK015)
CHAPTER 1 – MATTER 1
1.1 Atoms and molecules
1.2 Mole concept
1.3 Stoichiometry
CHAPTER 2 – ATOMIC STRUCTURE
2.1 Bohr’s atomic model
2.2 Quantum mechanical model
2.3 Electronic configuration
CHAPTER 3 – PERIODIC TABLE
3.1 Classification of elements
3.2 Periodicity
CHAPTER 4 – CHEMICAL BONDING
4.1 Lewis structure
4.2 Molecular shape and polarity
4.3 Orbital overlap and hybridization
4.4 Intermolecular forces
4.5 Metallic bond
CHAPTER 5 – STATES OF MATTER
5.1 Gas
5.2 Liquid
5.3 Solid
5.4 Phase diagram
CHAPTER 6 – CHEMICAL EQUILIBRIUM
6.1 Dynamic equilibrium
6.2 Equilibrium constants
6.3 Le Chatelier’s principle
CHAPTER 7 – IONIC EQUILIBRIA
7.1 Acids and bases
7.2 Acid-base titrations
7.3 Solubility equilibria
MARDIANA AMIR MIZAN
CHEMISTRY (SK015)
TABLE OF RELATIVE ATOMIC MASSES
Element Symbol Proton number Relative atomic mass
Silver Ag 47 107.9
Aluminum Al 13 27.0
Ar 18 40.0
Argon As 33 74.9
Arsenic Au 79 197.00
B 5 10.8
Gold Ba 56 137.3
Boron Be 4 9.0
Barium Bi 83 209.0
Beryllium Br 35 79.9
Bismuth C 6 12.0
Bromine Ca 20 40.1
Carbon Cd 48 112.4
Calcium Ce 58 140.1
Cadmium Cl 17 35.5
Cerium Co 27 58.9
Chlorine Cr 24 52.0
Cobalt Cs 55 132.9
Chromium Cu 29 63.6
Cesium F 9 19.0
Copper Fe 26 55.9
Fluorine H 1 1.0
Iron He 2 4.0
Hydrogen Hg 80 200.6
Helium I 53 126.9
Mercury K 19 39.1
Iodine Kr 36 83.8
Potassium Li 3 6.9
Krypton Mg 12 24.3
Lithium Mn 25 54.9
Magnesium N 7 14.0
Manganese Na 11 23.0
Nitrogen Ne 10 20.2
Sodium Ni 28 58.7
Neon O 8 16.0
Nickel P 15 31.0
Oxygen Pa 91 231.0
Phosphorus Pb 82 207.2
Protactinium Pt 78 195.1
Lead Ra 88 226.0
Platinum Rb 37 85.5
Radium Rn 86 222.0
Rubidium S 16 32.1
Radon Sb 51 121.8
Sulphur Sc 21 45.0
Stibium Se 34 79.0
Scandium Si 14 28.1
Selenium Sn 50 118.7
Silicon Sr 38 87.6
U 92 238.0
Tin W 74 183.9
Strontium Zn 30 65.4
Uranium
Tungsten 2
Zinc
MARDIANA AMIR MIZAN
CHEMISTRY (SK015)
LIST OF SELECTED CONSTANT VALUES
Ionisation constant for water at 25C Kw = 1.00 10−14 mol2 dm−6
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 = 6.021023 mol−1
Molar of gases constant
F = 9.65 x 104 C mol−1
Density of water at 25C
Freezing point of water h = 6.625610−34 J s
Vapour pressure of water at 25C
RH = 1.097 107 m−1
= 2.18 10−18 J
R = 8.314 J K−1 mol-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 1liter = 1 dm3
ENERGY 1mL = 1 cm3
PRESSURE 1J = 1 kg m2 s−2
OTHERS = 1Nm
1 calorie = 1 107 erg
1eV = 4.184 J
= 1.602 x 10-19 J
1 atm = 760 mmHg
= 760 torr
= 101.325 kPa
= 101 325 N m-2
1 faraday (F) = 96 500 coulomb
1 newton (N) = 1 kg m s−2
MARDIANA AMIR MIZAN 3
CHEMISTRY (SK015)
TUTORIAL 1.1 : ATOMS AND MOLECULES
Topic 1.1 : Atoms and Molecules
At the end of this topic, students should be able to :
a) Describe proton, electron and neutron in terms of the relative mass and relative charge.
b) Define proton number (Z), nucleon number (A) and isotope.
c) Write isotope notation.
d) Define relative atomic mass, Ar and relative molecular mass, Mr based on the C-12 scale.
e) Calculate the average atomic mass of an element given the relative abundance of isotopes or a
mass spectrum.
1 (a) Define :
i. proton number
ii. nucleon number
iii. isotopes
(b) Write the isotope notation for each of the following species:
Species Number of Notation
neutrons
Si - 28 protons electrons
Cl – 37 14 14 14
Na-23 17 20 18
11 12 10
2 (a) Define:
i. relative atomic mass
ii. relative molecular mass
(b) An atom of X is twice as heavy as one carbon-12 atom. Calculate the relative atomic mass
of element X. 24.00
3 The following is the mass spectrum of zirconium. Calculate the average atomic mass of zirconium.
52
%
intensity
14 13
9 12
90 91 92 93 94 mass(u)
MARDIANA AMIR MIZAN 91.27 amu
4
CHEMISTRY (SK015)
TUTORIAL 1.1 : ATOMS AND MOLECULES
4 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.
55.8514
5 The element Mg consists of three isotopes which are 24Mg, 25Mg and 26Mg. The relative atomic
mass of Mg is 24.3.
(a) Which is the most abundant isotope of Mg?
(b) State the number of protons and the number of neutrons of the isotope.
6 Chlorine isotopes occur naturally as 35Cl and 37Cl. The abundance ratio of these two isotopes is
35 Cl = 3.127. Based on the scale of carbon-12, the relative mass of 35Cl and 37Cl are 34.9689 and
37 Cl
36.9659 respectively. Calculate the relative atomic mass of chlorine.
35.4526
7 The isotopes of Ag occur naturally as 107Ag and 109Ag with their relative isotopic masses of
106.906 and 108.868 respectively. If the average atomic mass of Ag is 107.868, what would be
the percentage abundance of these two isotopes?
abundance of 107Ag = 50.97%
abundance of 108Ag = 49.03%
MARDIANA AMIR MIZAN 5
CHEMISTRY (SK015)
TUTORIAL 1.2 : MOLE CONCEPT
Topic 1.2 : Mole Concept
At the end of this topic, students should be able to :
a) Define mole in terms of mass of carbon-12 and Avogadro’s constant (NA)
b) Interconvert between moles, mass, number of particles, mola volume of gas at s.t.p and room
temperature.
c) Define the terms empirical and molecular formulae.
d) Determine empirical and molecular formulae from mass composition or combustion data.
e) Define and perform calculations for each of the following concentration measurements :
i. Molarity (M)
ii. Molality (m)
iii. Mole fraction (X)
iv. Percentage by mass (%w/w)
v. Percentage by volume (%v/v)
1 Calculate the relative molecular mass of each compound: 310.3
(a) calcium phosphate, Ca3(PO4)2 249.7
(b) hydrated copper(II) sulphate, CuSO4.5H2O
2 The relative atomic mass of element Yis 32.0. Calculate the mass of an atom of Y in gram.
5.3110-23 g
3 Calculate the number of moles for:
(a) 6.02×102 3 atoms of hydrogen. 1 mol H atoms
(b) 3.01×1021 molecules of ozone. 5.0010-3mol O3
(c) 24.5 g of hydrogen sulphate. 0.250 mol H2SO4
(d) 10.0 L O2 gas at STP. 0.446 mol O2
4 Determine:
(a) the number of atoms in 5.0 g of silver, Ag. 2.8 1022 Ag atoms
(b) the number of molecules in 25 g of methane, CH4. 9.41023 CH4 molecules
(c) the number of carbon atoms in 0.50mol of ethane, C2H4. 6.021023 C atoms
(d) the mass of carbon atoms in 1.08 g of quinine,C20H24N2O2 0.799 g of C
MARDIANA AMIR MIZAN 6
CHEMISTRY (SK015)
5 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 mLat STP. Determine the molar mass
and the molecular formula of the hydrocarbon. 56 g mol–1, C4H8
6 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.
CH2, C6H12
7 Define the terms:
(a) molarity.
(b) molality.
(c) percentage by mass.
8 The density of 10.5 molal NaOH solution is 1.33 g mL– 1 at 20C. Calculate: 0.159
29.6%
(a) the mole fraction of NaOH. 9.83 M
(b) the percentage by mass of NaOH.
(c) the molarity of the solution.
9 The density of 95% by mass of sulphuric acid, H2SO4is 1.84 g mL-1. Calculate
(a) the molarity of H2SO4. 17.8 M
(b) the volume of the acid needed to prepare 1.0 L of 0.080 M solution. 4.49 10–3 L
10 Hydrogen peroxide, H2O2, decomposes to produce water and oxygen as in the following equation:
2 H2 O2(l) ⎯→ 2H2O(l) + O2(g)
What would be the volume of O2 released from the decomposition of 250mL of 2.0 M H2O2 at
STP? 5.6 L
11 Ammonia reacts with oxygen according to the equation:
4NH3 (g) + 5O2 (g) ⎯→ 4NO(g) + 6H2O (l)
If 72.0 mL sample of NH3 gas is allowed to react with excess oxygen at room temperature, 25C
and 1 atm, calculate the number of molecules of water produced.
2.711021 H2O molecules
MARDIANA AMIR MIZAN 7
CHEMISTRY (SK015)
12 The complete combustion of copper(I) sulphide is according to the following equation:
2Cu2S(s) + 3O2(g)⎯→2Cu2O(s) + 2SO2(g)
If the mass of Cu2S in the mixture is 14.0 g, calculate 7.94 1022 O2 molecules
(a) the number of molecules of oxygen gas reacted. 5.63 g
(b) the mass of SO2 gas produced. 1.97 L
(c) the volume of SO2 gas at STP.
13 A sample of 0.40 g of anhydrous NaOH was dissolved in water and made up to 100 mL of solution.
25 mL of the solution was then neutralised with dilute H2SO4.
(a) Find the molarity of the NaOH solution. 0.1M
(b) How many moles of H2SO4 were required for the neutralisation?
0.00125 mol H2SO4
MARDIANA AMIR MIZAN 8
CHEMISTRY (SK015)
TUTORIAL 1.3 : STOICHIOMETRY
Topic 1.3 : Stoichiometry
At the end of this topic, students should be able to :
a) Determine the oxidation number of an element in a chemical formula.
b) Write and balance:
i.chemical equation by inspection method
ii.redox equation by ion-electron method
c) Define limiting reactant and percentage yield.
d) Perform stoichiometric calculations using mole concept including limiting reactant and
percentage yield.
1 Determine the oxidation number of the underlined elements in the following
compounds:
(a) NO2 +4
(b) KMnO4 +7
(c) Cr2O72- +6
(d) IO3- +5
2 Balance each of the following equations:
(a) NaOH(aq) + FeCl3(s) ⎯→ Fe(OH)3(s) + NaCl(aq)
(b) Fe2O3(s) + HCl(aq) ⎯→ FeCl3(aq) + H2O(l)
(c) Cr(OH)3(aq) + IO3-(aq) ⎯→ CrO32-(aq) + I-(aq) (acidic medium)
(d) Cl2(aq) ⎯→ ClO4–(aq) + Cl–(aq) (basic medium)
3 A reaction between 70 g of copper(II) oxide and 50 mL of 2.0 M nitric acid produces copper(II)
nitrate, Cu(NO3)2 and water.
(a) Define limiting reactant.
(b) Write the balanced chemical equation for the above reaction.
(c) Determine the limiting reactant.
(d) Calculate the mass of excess reactant after the reaction. 66.02g
(e) Determine the percentage yield if the actual mass of copper(II) nitrate obtained from the
reaction is 8.5 g. 90.63%
4 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 mLof 0.020 M KMnO4 is requiredfor the titration to reach the
equivalence point.
(a) Write the balanced equation for the titration.
(b) Calculate the percentage of iron in the sample. 33.49 %
MARDIANA AMIR MIZAN 9
CHEMISTRY (SK015)
TUTORIAL 2.1 : BOHR’S ATOMIC MODEL
Topic 2.1 : Bohr’s Atomic Model
At the end of this topic, student should be able to :
(a) Describe Bohr’s atomic model.
(b) Explain the existence of energy levels in an atom
(c) Calculate the energy of an electron using: En = RH( 1 2), RH = 2.18 x 10-18 J
(d) Describe the formation of line spectrum of hydrogen atom
(e) Calculate the energy change of an electron during transition
(f) Calculate the photon of energy emitted by an electron that produces a particular wavelength
during transition.
(g) Perform calculations involving the Rydberg equation for Lyman, Balmer, Paschen, Brackett
and Pfund series.
(h) Calculate the ionization energy of hydrogen atom from Lyman series.
(i) State the weaknesses of Bohr’s atomic model.
(j) State the dual nature of electron using de Broglie’s postulate and Heisenberg’s uncertainty
principle.
1 Describe Bohr’s atomic postulates.
2 (a) How is the second line of Brackett series produced?
(b) The first line of Balmer series has a wavelength of 656.3 nm. Calculate its frequency.
4.5711014 s-1
(c) Calculate the wavelength (in nm), energy and the frequency of the fourth line in Lyman
series. 94.95 nm,3.16 x 1015 s-1 , 2.0946 x 10-18J
3 With reference to the table below: Wavelength /nm
656.3
Colour of line spectrum 486.3
Red 434.2
Green
Blue
(a) Calculate the energy emitted in the formation of red line and green line.
-4.09010-19 J, -3.03110-19 J
(b) Sketch the energy level diagram to show the transitions of electron that produced the red,
green and blue lines.
MARDIANA AMIR MIZAN 10
CHEMISTRY (SK015)
TUTORIAL 2.1 : BOHR’S ATOMIC MODEL
4 Figure below shows the Lyman series of hydrogen emission spectrum.
i) Label the low and high energy ends
ii) Draw the electron transition of lines P, Q and R on the energy level diagram of the hydrogen
atom.
iii) Calculate the energy corresponding to line Q. - 2.04 x 10 -18 J
5 The difference in energy between the second and fourth energy levels of a hydrogen atom is
4.0910-19 J. Calculate;
a) The wavelength and frequency of the photon emitted when an electron makes a
transition from the fourth to the second energy level.
4.86 x 10-7 m , 6.17 x 1014 s-1
b) Energy emitted by 1 mole of electrons for the above transition. 2.46x102 kJ mol-1
6
The Lyman series of the spectrum of hydrogen is shown above. Calculate the ionization energy
of hydrogen from the spectrum. 2.1819 x 10-18 J , 1313kJmol-1
7 (a) Give two weaknesses of Bohr’s atomic model.
(b) State de Broglie’spostulate on the dual nature of electron and Heisenberg’s uncertainty
principle.
MARDIANA AMIR MIZAN 11
CHEMISTRY (SK015)
TUTORIAL 2.2 : QUANTUM MECHANICAL MODEL
Topic 2.2 : Quantum Mechanical Model
At the end of this topic, student should be able to :
(a) Define the term orbital.
(b) State all the 4 quantum number of an electron in an orbital n,l,m,s.
(c) Sketch the 3-D shapes of s,p and d orbitals.
1 Give one set of quantum numbers for an electron in 2s, 4p and 3d orbitals.
2 (a) Which of the following orbitals are allowed?
(i) 2d (ii) 7s (iii) 3f (iv) 4p
(b) For each set of the quantum numbers; b C
set a 3 5
n2 2 1
l0
(i) State name of the orbitals.
(ii) How many orbitals in each set.
(iii) Determine the maximum number of electrons that can occupy each orbital.
3 Which of the following quantum numbers (n, l, m, s), 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, +½)
4 Calcium is the chemical element with symbol Ca and atomic number 20. Calcium is a soft gray
alkaline earth metal, and is the fifth-most-abundant element by mass in the Earth's crust.
a) Find the number of electrons with
i) principal quantum number, n, equals to 4
ii) angular momentum quantum number,l, equals to 1
iii) magnetic quantum number, m, equals to 0
iv) spin quantum number, s, equals to -1/2
b) Draw the shape of the orbitals when n= 3.
MARDIANA AMIR MIZAN 12
CHEMISTRY (SK015)
TUTORIAL 2.3 : ELECTRONIC CONFIGURATION
Topic 2.3 : Electronic Configuration
At the end of this topic, student should be able to :
a) State Aufbau principle, Hund’s rule and Pauli exclusion principle.
b) Apply the rules in (a) to fill electron into atomic orbitals
c) Write the electronic configuration of atoms and monoatomic ions using spdf notation
d) Explain the anomalous electronic configurations of chromium and copper.
1 Define:
(a) Aufbau’s principle.
(b) Hund’s rule
(c) Pauli Exclusion Principle.
2 Write the electronic configuration for the following atoms/ions using
(a) spdf notation and orbital diagram
i. Na ii. Na+ iii. Cl iv Cl−
(b) i) Explain the anomalous electronic configuration in chromium and copper.
ii) Give the possible value of n, l, m of third electron in actual configuration of copper.
MARDIANA AMIR MIZAN 13
CHEMISTRY (SK015)
TUTORIAL 3.1 : CLASSIFICATION OF ELEMENTS
Topic 3.1 : Classification of Elements
At the end of this topic, students should be able to :
a) Describe period, group and block (s,p,d,f).
b) Specify the position of metals, metalloids and non-metals in the periodic table.
c) Deduce the position of elements in the periodic table from its electronics configuration.
1. State the period, group and block for each element below:
A : 1s2 2s2 2p6 3s2 3p6
B : 1s2 2s2 2p6 3s2 3p6 4s2
C : 1s2 2s2 2p6 3s2 3p6 3d3 4s2
D : 1s2 2s2 2p6 3s2 3p6 4s2 3d7
2. The table below shows the group and period of the elements P, Q, R and S.
Element Period Group
P 2 16
Q 2 18
R 3 1
S 4 17
(a) Write the electronic configuration for all of the elements.
(b) Write the molecular formulae of the compound formed when P react with R.
3 The figure below shows the location of elements A, B, C, D and Ein periodic table.
A FC D
B
E
(a) Which of the elements are metals, non-metals and metalloids ?
(b) Arrange the elements B, C, D and E in ascending order of atomic radius. Explain.
MARDIANA AMIR MIZAN 14
CHEMISTRY (SK015)
TUTORIAL 3.2 : PERIODICITY
Topic 3.2 : Periodicity
At the end of this topic, students should be able to :
a) Explain the variation in atomic radii.
b) Compare the atomic radius of an element and its corresponding ionic radius.
c) Define the term isoelectronic.
d) Explain the radius of isoelectronic species.
e) Explain the variation in the ionic radii across period 2 & period 3.
f) Define the first and second ionisation energies.
g) Explain the variations in the first ionisation energy across a period and down a group.
h) Explain the increase in the successive ionisation energies of an element.
i) Deduce the electronic configuration of an element and its position in the periodic table
based on successive ionisation energy data.
j) Define electronegativity.
k) Explain the variation in electronegativity of elements.
l) Describe the periodicity of elements across period 3 and down Group 1 and 17 for the
following physical properties - Metallic character , Melting point , Boiling point
m) Explain the acid-base character of oxides of elements in Period 3.
1 Removal and addition of electron results in change of atomic radii.
Species Na Na+ Cl Cl-
0.095 0.099 0.181
Radius/nm 0.156
Explain the difference in radius between the ions and their respective neutral atoms.
2 Which ion has the smaller radius? Explain
(a) Mg2+ or Na+
(b) Na+ or O2-
(c) Cl- or Na+
3 The table below shows data for successive ionization energy of three elements X, Y and Z.
Elements Ionisation Energy (kJ/mol)
X First Second Third Fourth Fifth
Y 1086 38000
Z 786 2350 4620 6200 16000
738 13600
1580 3230 4360
1450 7730 10500
(a) Predict the group of X in the periodic table. Give your reason.
(b) Write the valence electronic configuration of Z.
MARDIANA AMIR MIZAN 15
CHEMISTRY (SK015)
4 The graph below shows the variation of ionisation energies for element Z.
Log ionisation energy
number of electrons removed
(a) State the group and period for element Z.
(b) Why a higher energy is needed to remove the fourth electron?
(c) State the oxidation number of element Z.
5 Which element has the higher first ionization energy (IE1)?
(a) Mg or Na
(b) O or S
(c) N or O
6 (a) Determine which of the element has higher boiling point. Explain.
i) Na and Al iii) Li and Na
ii) Si and S iv) F and Cl
(b) Between Mg and Al, which atom has a higher metallic character? Explain your answer.
7 TABLE 1 shows the position of elements P, Q, R, S and T in the periodic table.
Element Period Group
P 2 14
Q 3 1
R 3 15
S 3 16
T 4 1
TABLE 1
(a) State the proton number for element Q.
(b) Which of the elements has the greatest metallic character?
MARDIANA AMIR MIZAN 16
CHEMISTRY (SK015)
(c) Arrange the elements in TABLE 1 in the ascending order of atomic size.Explain.
(d) Give ions that are isoelectronic with Argon.
(e) Explain why first ionisation energy R is higher than S.
(f) List down the elements in TABLE 1 that will form acidic oxides.
8 FIGURE 2 show part of a periodic table. The positions of nine elements are indicated byletters
not representing their usual symbols. Answer the following questions based on this figure.
AB
C DE F
GH I
FIGURE 2
Identify the element which
(a) has the largest ionic radius.
(b) has the smallest atomic radius.
(c) has the lowest first ionisation energy.
(d) has the highest electronegativity.
(e) forms amphoteric salt with oxygen.
MARDIANA AMIR MIZAN 17
CHEMISTRY (SK015)
TUTORIAL 4.1 : LEWIS STRUCTURE
Topic 4.1 : Lewis Structure
At the end of this topic, students should be able to :
a) write the Lewis structure for an atom.
b) state the octet rule and describe how atoms obtain the octet configuration.
c) describe the formation of the following bonds using Lewis symbol :
i. Ionic or electrovalent bond.
ii. Covalent bond
iii. Dative or coordinate bond.
d) draw Lewis structure of covalent species 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: incomplete octet, expanded octet and odd number
electrons.
h) explain the concept of resonance using appropriate examples.
1 Define Lewis symbol.
2 Complete the table below :
Group 1 2 13 14 15 16 17 18
No. of valence electron
Lewis dot symbol Na Mg Al Si P S Cl Ne
3 For each of the following elements, write the electronic configuration for its most stable ion and
state the type of stability:
(a) 13Al (b) 9F (c) 26Fe
(d) 7N (e) 16S (f) 30Zn
4 (a) How is an ionic bond formed?
(b) By using Lewis symbol, show the formation of:
i. MgBr2
ii. Al2O3
iii. BaO
5 (a) Using elements sodium and chlorine as examples, explain the differenc between an ionic
bond and a covalent bond.
(b) Use Lewis symbols to show the sharing of electrons between hydrogen and fluorine 18
atoms. Label the electron pairs as bonding pair(s) and lone pair(s).
MARDIANA AMIR MIZAN
CHEMISTRY (SK015)
6 Describe the formation of dative bond by using NH4+ and NH3BF3 as
examples.
7 Draw the Lewis structure for each of the following molecules and state the unusual features
(exceptional octet) of the structure.
(a) ClF3
(b) BCl3
(c) SeF4
(d) NO
(e) AlBr3
8 Draw the Lewis structure for POCl3 and determine the formal charge for each
atom in POCl3.
9. Draw the resonance structures for NO3- and SO3.
10. 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.
(b) Determine the most plausible structure. Explain.
MARDIANA AMIR MIZAN 19
CHEMISTRY (SK015)
TUTORIAL 4.2 : MOLECULAR SHAPE AND POLARITY
Topic 4.2 : Molecular Shape and Polarity
At the end of this topic, students should be able to :
a) Explain valence shell electron pair repulsion theory (VSEPR)
b) Draw the basic molecular shapes: linear, trigonal planar, tetrahedral, trigonal bipyramidal and
octahedral (include bond angles)
c) Predict and explain 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.
1 (a) State the valence shell electron pair repulsion theory
(b) Draw and state the geometry of the species in which around the central atom there are,
i. Five single bonds, one lone pair electrons
ii. Four single bonds, two lone pair electrons.
iii. Five single bonds.
iv. Three single bonds, two lone pair electrons.
v. Two single bonds, two lone pair electrons.
vi. Two single bonds, one lone pair electrons.
2 Use VSEPR theory to predict the molecular geometries of the following species :
(a) O3
(b) PBr5
(c) NO2+
(d) SnCl3-
(e) XeF3+
(f) XeF5+
(g) SbF6-
3 The bond angle in the CH4, NH3 and H2O molecules are 109.5o, 107o and 104.5ᵒ respectively.
Explain according to VSEPR Theory.
4 (a) Define dipole moment. What is the formula of dipole moment?
(b) How is it possible for a molecule to have a polar bond and yet to be non-polar?
(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
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CHEMISTRY (SK015)
TUTORIAL 4.3 : ORBITAL OVERLAP AND HYBRIDIZATION
Topic 4.3 : Orbital Overlap and Hybridization
At the end of this topic, students should be able to :
a) Draw and describe the formation of sigma(σ) and pi(п) bonds from overlapping of orbitals.
b) Draw and explain the formation of hybrid orbitals of a central atom: sp, sp2, sp3, sp3d, sp3d2
using appropriate examples.
c) Draw orbitals overlap and label sigma(σ) and pi(п) bond of a molecule.
1 (a) Describe the formation of sigma(σ) and pi(Π) bonds.
(b) Use valence bond theory to explain the bonding in HF and show how the bond is formed
by the overlapping of atomic orbitals.
2 (a) What is meant by hybridisation of atomic orbitals ?
(b) Indicate the type of hybridisation of carbon and nitrogen atoms in the molecule below
(Hint : you have to draw the Lewis structure of the species to answer this question)
NH2-C(NH)-NH-CH2-CN
(c) State the type of the hybridization of the underlined atoms in the following compounds
i. CH3CH3
ii. CH3CH=CH2
iii. BF3
iv. CCl4
3 Describe the hybridisation process and determine the geometry of the following species.
(a) BeH2
(b) BH3
(c) PF5
(d) SeF4
(e) SCl6
(f) ICI2-
(g) ICI4+
MARDIANA AMIR MIZAN 21
CHEMISTRY (SK015)
TUTORIAL 4.4 : INTERMOLECULAR FORCES
Topic 4.4 : Intermolecular Forces
At the end of the lesson, students should be able to :
a) Describe intermolecular forces van der Waals forces (dipole-dipole interactions or permanent
dipole & London dispersion forces) and hydrogen bonding.
b) Explain factors that influence the strength of van der Waals forces and hydrogen bond.
c) Explain the effects of hydrogen bonding on boiling point, solubility, density of water compared
to ice.
1 (a) What are dipole-dipole forces?
(b) How do polar molecules attract one another?
(c) Explain how London forces arise between non-polar molecules?
2 (a) Name the type of van der Waals forces that exist between molecules for each of the
following species:
i. CH4
ii. SO2
iii. I2
iv. H2S
v. C6H6
3 (a) State the factors that influence the strength of van der Waals forces?
(b) Which molecule has a higher boiling point. Br2 or ICl? Explain.
4 (a) Describe hydrogen bond.
(b) Ethanol, C2H5OH and methyl ether CH3OCH3 have the same molar mass. Which has a
higher boiling point?
(c) Ethanol (C2H5OH, molar mass 46) boils at 78 oC, but water ( H2O, molar mass 18) boils
at 100 oC. Explain.
5 Arrange the following compounds in order of increasing solubility in water.
CH3CH2OH, CH3CH3 and NH3 .
6 Explain why ice is less dense than water.
MARDIANA AMIR MIZAN 22
CHEMISTRY (SK015)
TUTORIAL 4.5 : METALLIC BOND
Topic 4.5 : Metallic Bond
At the end of the lesson, students should be able to :
a) Explain the formation of metallic bonding by using electron sea model.
b) Relate metallic bond to the properties of metal :
i. malleability
ii. ductility
iii. electrical conductivity
iv. thermal conductivity
c) Explain the factors that affect the strength of metallic bond
d) Relate the strength of metallic bond to boiling point
1 (a) What is meant by a metallic bond.
(b) Describe the formation of metallic bond in aluminium by using electron sea model.
2 Explain the difference in electrical conductivity of magnesium and sulphur.
3 The melting points of aluminium, Al, and aluminium chloride, AlCl3 are 660 °C and 192 °C
respectively. Explain the difference in boiling points in terms of chemical bonding.
4 (a) Explain why the melting point of aluminium is higher then that of sodium.
(b) Explain why the boiling point of calcium higher than that of potassium.
MARDIANA AMIR MIZAN 23
CHEMISTRY (SK015)
TUTORIAL 5.1 : GAS
Topic 5.1 : Gas
At the end of this topic, students should be able to :
a) Explain qualitatively the basic assumption of the kinetic molecular theory of gases for an ideal
gas.
b) Define gas laws :
i. Boyle’s law
ii. Charles’s law
iii. Avogadro’s law
c) Sketch and interpret the graphs of Boyle’s and Charles’s law.
d) Perform calculations involving gas laws and ideal gas equation.
e) Define and perform calculation using Dalton’s law.
f) Explain the ideal and non-ideal behaviours of gases in terms of intermolecular forces and
molecular volume.
g) Explain the condition at which real gases approach the ideal behaviour.
1 State the assumptions of the ideal gas behaviour.
2 (a) State Boyle’s law.
(b) 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? 0.40 atm
3 (a) State Charles’s law.
(b) A sample of gas occupies 100.0 mL at 25C. What volume would the gas occupy at 32C
if the pressure remains constant? 102 mL
4 (a) Derive the ideal gas equation from the ideal gas laws.
(b) Derive an equation that relates the density of gas, , to its pressure, P, from the ideal gas
equation.
(c) Gas C with a mass of 8 g at 0°C and 0.87 atm occupies a volume which is equal to a volume
of 11g CO2 at STP. Calculate the density and the molar mass of the gas.
36.84 gmol-1
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 31C, 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.
69.33
6 (a) State Dalton’s law of partial pressure.
(b) 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 total pressure of the gases at 25oC? 2.36 atm
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CHEMISTRY (SK015)
7 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
8 Describe the differences between the ideal gas equation and the van der Waals equation
MARDIANA AMIR MIZAN 25
CHEMISTRY (SK015)
TUTORIAL 5.2 : LIQUID
Topic 5.2 : Liquid
At the end of this topic, students should be able to :
a) Explain the properties of liquid: shape, volume, surface tension, viscosity, compressibility and
diffusion.
b) Explain vaporisation and condensation processes based on kinetic molecular theory and
intermolecular forces.
c) Define vapour pressure and boiling point.
d) Explain the relationship between :
i. intermolecular forces and vapour pressure.
ii. vapour pressure and boiling point.
1 (a) Describe properties of liquid.
(b) Based on kinetic-molecular theory, explain vaporisation and condensation process.
2 (a) Define the terms vapour pressure and boiling point. Boiling point/K
(b) Based on Table 2, 231
Table 2 351
Compound 313
Propane, CH3CH2CH3
Ethanol, CH3CH2OH
Dimethyl ether, CH3OCH3
i. Arrange the compounds in order of increasing strength of intermolecular forces.
Explain your answer.
ii. Which of the following compound will have the highest vapour pressure? Explain.
MARDIANA AMIR MIZAN 26
CHEMISTRY (SK015)
TUTORIAL 5.3 : SOLID
Topic 5.3 : Solid
At the end of this topic, students should be able to :
a) Explain the fixed-shape of a solid.
b) Explain the process of :
i. Freezing (solidification)
ii. Melting (fusion)
iii. Sublimation
iv. Deposition
1 Differentiate between amorphous and crystalline solids.
2 Give an example for each of the following crystalline solids:
(a) metallic
(b) ionic
(c) molecular covalent
(d) giant covalent
3 Fill in the blanks in the following table to give a brief comparison between a solid and a liquid.
Property Solid Liquid
Shape
Volume
Compressibility
Kinetic energy of particles
4 When a solid is heated, it undergoes phase changes. Describe the processes.
5 Describe the melting process in term of the molecular kinetic theory.
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CHEMISTRY (SK015)
TUTORIAL 5.4 : PHASE DIAGRAM
Topic 5.4 : Phase Diagram
At the end of this topic, students should be able to :
a) Define phase, triple point and critical point.
b) Sketch and differentiate the phase diagram of H2O and CO2.
c) Identify triple and critical point on the phase diagram.
d) Explain the anomalous behaviour of H2O.
e) Describe the changes in phase with respect to
i. Temperature (at constant pressure)
ii. Pressure (at constant temperature)
1 c) Define phase and component.
2 State the number of components and phases for each of the following system.
(a) A mixture of ethanol and water.
(b) A solution of benzene and water.
(c) A mixture of O2, N2 and CO2 gases.
(d) HNO3 (aq) in equilibrium with its vapour.
3 (a) Define critical point and triple point.
(b) TABLE 1 shows data phase equilibrium diagram for carbon dioxide.
TABLE 1
Triple point Pressure (atm) Temperature (°C)
Critical point 5.2 -57
73 31
i. Sketch a phase diagram for carbon dioxide.
ii. By referring to the phase diagram, explain why solid carbon dioxide at room
temperature and pressure does not melt but sublimes.
iii. Explain the phase changes that could possibly occur when a sample of carbon dioxide
in a closed vessel under 1 atm pressure and a temperature of -78°C is pressurised
isothermally to 10 atm, followed by isobaric heating to 20°C.
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CHEMISTRY (SK015)
4 FIGURE 1 below shows the phase diagram of water.
P
CA
◘E ◘F
D
B
Temperature
FIGURE 1
(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. (hint: draw and label new points in your diagram to explain your answer)
(c) Name the point at which the BC line intersects 1 atm line.
(d) The BC line has a negative slope. Explain.
6 TABLE 2 shows phase equilibrium data for substance D.
Sublimation point Pressure (atm) Temperature (°C)
Triple point 2.5 -80
Critical point 5.0 -60
Melting point 1 40.0 30
Melting point 2 15.0 -30
Boiling point 1 35.0 -15
Boiling point 2 20.0 10
30.0 25
Plot and label a phase diagram for D. Does D sublime at 20 atm ? Explain your answer.
MARDIANA AMIR MIZAN 29
CHEMISTRY (SK015)
TUTORIAL 6.1 : DYNAMIC EQUILIBRIUM
Topic 6.1 : Dynamic Equilibrium
At the end of this subtopic, students should be able to :
a) Explain :
i) reversible reaction
ii) dynamic equilibrium
iii) law of mass action
b) State characteristics of a system in equilibrium.
c) Explain the change of concentrations of reactants and products based on the curve of
concentration against time for reversible reaction.
d) Define homogeneous and heterogeneous equilibria.
e) Write expressions for equilibrium constants in term of concentration (Kc) and partial pressure
(Kp) for homogeneous and heterogeneous systems.
1 d) (a) Define the following terms: ii) law of mass action (equilibrium law)
(b) i) reversible reaction. iv) heterogeneous equilibrium
iii) homogeneous equilibrium
State two characteristics of a dynamic equilibrium reaction.
2 For the reversible reaction below :
AB
Sketch a graph of the variation of the concentration of A and B with time untill the system has
achieved equilibrium.Explain the shape of your graph.
3 (a) Determine whether the following reactions are homogeneous or heterogeneous.
i. 2PCl3(g) + O2(g) 2POCl3(g)
ii. FeO(s) + CO(g) Fe(s) + CO2(g)
iii. Ag+ (aq) + Fe2+(aq) Ag(s) + Fe3+(aq)
iv. 3Fe(s) + 4H2O(l) Fe3O4(s) + 4H2(g)
(b) Write the equilibrium law for the above reactions in terms of concentration, Kc, and partial
pressure, Kp.
MARDIANA AMIR MIZAN 30
CHEMISTRY (SK015)
TUTORIAL 6.2 : EQUILIBRIUM CONSTANTS
Topic 6.2: Equilibrium Constants
At the end of this subtopic, students should be able to :
a) Derive and use the equation, Kp = Kc (RT)Δn
b) Calculate Kc, Kp or the quantities of the species present at equilibrium
c) Define and determine degree of dissociation, α.
d) Deduce the expression for the reaction quotient, Q and predict the direction of nett reaction by
comparing the values of Q and K
1 At equilibrium, there are 2.50 moles of SO2 , 1.35×10-5 moles of O2 and 8.70 moles of SO3 present
in a 12.0 L flask. Calculate the equilibrium constant, Kc.
2SO2(g) + O2(g) 2SO3(g) 1.077 x 107
2 The partial pressures of PCl3, O2 and POCl3 at equilibrium are 0.0498, 0.0247 and 0.9967 atm
respectively. Calculate the equilibrium constant, KP, for the reaction.
2PCl3(g) + O2(g) 2POCl3(g) 1.622 x 104
3 The following reaction achieved equilibrium when the partial pressure of bromine gas is 0.600
atm.
FeBr3(s) FeBr2(s) + ½ Br2(g)
(a) Determine KP. 0.775
(b) Derive KP = KC (RT)1/2.
4 At 25oC, H2O(l) reached equilibrium according to the equation,
H2O(l) H2O (g)
Calculate Kp and Kc for the reaction. 0.0313 ; 1.28x10-3
[vapour pressure of water at 25oC = 23.8 torr]
5 1.00 mol each of CO and Cl2 are introduced into an evacuated 1.75 L flask at 668K. At
equilibrium, the total pressure of the gaseous mixture is 32.4 atm. Calculate Kp.
CO(g) + Cl2(g) COCl2(g) 25.9
6 Ammonium hydrogen sulfide decomposes at 250oC according to the reaction.
NH4HS (s) NH3 (g) + H2S (g) Kp = 0.110
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 ; 0.33 atm
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CHEMISTRY (SK015)
7 Bromine gas is allowed to reach equilibrium at 1280oC according to the equation,
Br2 (g) 2Br (g) Kc = 1.10 × 10-3
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.
0.065 M ; 0.0084 M
8 (a) Define degree of dissociation
(b) 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, α of N2O4. 0.12
9 At 430 oC, 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 Chloromethane is formed by the reaction;
CH4(g) + Cl2(g) CH3Cl(g) + HCl(g) Kp = 1.6×104 at 1500 K
In the reaction mixture, the partial pressures of CH4, Cl2 , CH3Cl and HCl are 0.13 atm, 0.035
atm, 0.24 atm and 0.47 atm respectively. Determine whether more CH3Cl or CH4 is formed.
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TUTORIAL 6.3 : LE CHATELIER’S PRINCIPLE
Topic 6.3 : Le Chatelier’s Principle
At the end of this subtopic, students should be able to :
a) State Le Chatelier’s Principle
b) Explain the effect of the following factors on a system at equilibrium using Le Chatelier’s
Principle :
a) concentrations of reacting species
b) pressure/volume
c) addition of inert gas at constant volume and at constant pressure
d) temperature
e) catalyst
1 (a) State Le Chatelier’s principle.
(b) 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 of the reaction if :
i) the concentration of hydrogen gas is decreased.
ii) the temperature is lowered.
iii) a catalyst is added.
iv) C2H6 is removed from the system.
v) the volume of the container is increased.
vi) the pressure is increased.
vii) an inert gas is added at constant pressure.
viii) an inert gas is added at constant volume.
2 The Haber process is represented by this equation
N2(Ng2)( +g)3+H3H2(2(g)g ) 22NNHH3( 3g(g)) ΔH0 = - 92.2 kJ
Its feasibility depends on choosing conditions under which nitrogen and hydrogen reacts rapidly
to give a high yield of ammonia. At 25 0C and atmospheric pressure, the position of the
equilibrium favors the formation of NH3. Based on Le Chatelier’s principle, explain how the
product could be increased.
3 The decomposition of ammonium hydrogen sulfide is an endothermic process.
NH4HS(s) NH3(g) + H2S(g)
A 6.80 g sample of the solid is placed in an evacuated 4.0 L vessel at 24 °C. After equilibrium has
been established, the total pressure inside is 0.709 atm. Some solid NH4HS remains in the vessel.
What is the Kp for the reaction? Determine percentage of the solid has decomposed.
If the volume of the vessel were doubled at constant temperature, what would happen to the amount
of solid in the vessel? What happen to percentage of dissociation, σ ?
MARDIANA AMIR MIZAN 33
CHEMISTRY (SK015)
TUTORIAL 7.1 : ACIDS AND BASES
Topic 7.1 : Acids and Bases
At the end of this topic, students should be able to :
a) Define acid and base according to Arrhenius, Bronsted-Lowry and Lewis theories.
b) Deduce conjugate acid and conjugate base according to Bronsted Lowry theory.
c) Define strong acid and base, weak acid and base, pH and pOH.
d) Relate pH and pOH to the ionic product of water (Kw) at 25 oC.
e) Calculate the pH values of a strong acid and base.
f) Relate the strength of a weak acid and a weak base to the respective dissociation constants, Ka
and Kb.
g) Perform calculations involving pH, dissociation constant, initial concentration, equilibrium
concentration and the degree of dissociation (α).
h) Explain salt hydrolysis equations and classify the salts formed from the reaction between
i. Strong acid and strong base
ii. Strong acid and weak base
iii. Weak acid and strong base
i) Define buffer solution
j) Describe how a buffer solution controls its pH
1. (a) Define acid and base according to Bronsted-Lowry theory.
(b) Identify the conjugate acid-base pairs in each of the following reactions:
i) CH3COO− + HCN CH3COOH + CN−
ii) H2PO4− + NH3 HPO42− + NH4+
2 (a) i) Define pH.
ii) Calculate the pH of a 0.025 M sulphuric acid, H2SO4. 1.30
iii) Calculate the mass of NaOH needed to prepare 500.0 mL of solution with a pH of
10.00 ? 0.002g
3 (a) Write an expression for the dissociation constant, Ka of propanoic acid, CH3CH2COOH.
(b) Calculate the pH of a 0.35 mol L−1 solution of propanoic acid at 25oC. 2.66
[Ka = 1.3510−5 mol L-1]
4 (a) The pH of a 0.100 M solution of a weak acid, HA, is 2.85. What is the Ka of the acid?
2.02x10-5
(b) Calculate the percentage ionisation of 0.10 M hydrocyanic acid, HCN solution.
[Ka HCN = 5.0 10−10 mol L−1] 7.1x10-3 %
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5 What is the pH of a 0.10 M solution of phenylamine, C6H5NH2 ? 8.8
[Kb C6H5NH2 = 4.0 10−10 mol L−1]
6 TABLE 1 shows the base ionisation constant, Kb, for several selected compounds.
(a)
Compound Kb
C6H5NH2 3.810−10
1.710−6
N2H4 1.810−5
NH3 1.110−8
NH2OH
TABLE 1
i. Arrange the compounds in order of increasing strength of base.
ii. Give the structure of conjugate acid for each compound and arrange them in order
of increasing strength of acid.
(b) The percentage ionisation of 0.010 M NH3 solution was 4.2 % ionisation. Calculate Kb.
1.8x10-5
7 Write the salt hydrolysis equation for the following salts and classify them as acidic, basic or
neutral.
(a) NaCN
(b) N2H5Cl
8 (a) Calculate the pH of a solution containing 0.20 M CH3COOH and 0.30 M CH3COONa.
[Ka CH3COOH = 1.8 x 10-5] 4.9
(b) Calculate the pH of the 0.20 M CH3COOH solution if there is no salt present. 2.7
[Ka CH3COOH = 1.810−5]
(c) Explain the change in pH the solution in (a) when
i. a small amount of strong acid is added.
ii. a small amount of strong base is added.
9 Calculate the mass of ammonium chloride, NH4Cl needed to dissolve into 30 mL 0.15M ammonia
solution to form a buffer with a pH of 8.3.
[Kb NH3 = 1.810−5] 2.2g
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CHEMISTRY (SK015)
TUTORIAL 7.2 : ACID-BASE TITRATIONS
Topic 7.2 : Acid-base Titrations
At the end of this subtopic, students should be able to :
a) Describe titration process and distinguish between the end point and equivalence point
b) Sketch and interpret the variation of pH against titrant volume for titrations between
i. Strong acid-strong base
ii. Strong acid-weak base
iii. Weak acid-strong base
1 Calculate pH of the solution formed when: 1.18
(a) 150.00 mL of 0.200 M HNO3 is mixed with 75.00 mL of 0.200 M NaOH. 7
(b) 25.00 mL of 0.450 M H2SO4 is mixed with 25.00 mL of 0.900 M NaOH.
(c) 20.00 mL of 0.100 M HNO3 is mixed with 30.00 mL of 0.100 M NaOH 12.30
2 Calculate the pH of the solution if 50.00 mL of 0.20 M of NH3 is mixed with 30.00 mL of 0.15M
of HBr. [Kb of NH3 = 1.8 x 10-3] 9.34
3 Sketch a titration curve and give the suitable indicator for the following condition of titration:
a) 25.0 mL of 0.1 M HCl with 0.1 M NaOH
b) 25.0 mL of 1.0 M NH3 with 0.25 M HCl
c) 30.0 mL of 0.1 M CH3COOH with 0.05 M NaOH
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CHEMISTRY (SK015)
TUTORIAL 7.3 : SOLUBILITY EQUILIBRIA
Topic 7.3 : Solubility Equilibria
At the end of this subtopic, students should be able to :
a) Define solubility, molar solubility and solubility product (Ksp).
b) Calculate Ksp from concentrations of ions and vice versa.
c) Predict the possibility of precipitation of slightly soluble ionic compounds by comparing the
values of ion-product (Q) to Ksp.
d) Define and explain the common ion effect.
1 (a) Define solubility.
(b) The solubility of silver sulphate is 1.510-2 mol L−1. Calculate the solubility product of
the salt. 1.4x10-5
2 It was found experimentally that the Ksp of calcium sulphate is 2.410−4. Calculate the molar
solubility of calcium sulphate. 1.6x10-2 M
3 Will precipitate form if 200 mL of 0.004 M BaCl2 is added to 600 mL of 0.008 M K2SO4 ?
[Ksp BaSO4= 1.110−10]
4 Calculate the solubility of silver chromate, Ag2CrO4 at 25C in 1.3x10-4 M
(a) pure water. 2.12x10-5 M
(b) solution of 0.005 M K2CrO4 solution.
[Ksp Ag2CrO4 = 9.010−12]
Explain the difference in the values of solubility.
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HANDOUT 1
NAMES AND FORMULA OF SOME COMMON INORGANIC CATIONS
AND ANIONS
Cations Names Anions Names
H+ hydrogen ion fluoride
Li+ lithium ion F¯ chloride
Na+ sodium ion bromide
K+ potassium ion Cl¯ iodide
Be2+ beryllium ion oxide
Mg2+ magnesium ion Br¯ sulfide
Ca2+ calcium ion hydride
Ba2+ barium ion I¯ nitrate
Ag+ O2- perchlorate
Zn2+ silver ion S2- nitrite
Al3+ zinc ion chlorate
Cu+ aluminum ion H¯ chromate
Cu2+ copper (I) ion NO3¯ chlorite
Fe2+ copper (II) ion ClO4¯ dichromate
Pb2+ iron (II) ion NO2¯ hypochlorite
Fe3+ lead (II) ion lO3¯ cyanide
Sn2+ iron (III) ion CrO42- periodate
Co2+ tin (II) ion ClO2¯ permanganate
Sn4+ cobalt (II) ion Cr2O72- iodate
Co3+ tin (IV) ion ClO¯ hydroxide
Cr2+ cobalt (III) ion CN¯ peroxide
Cr3+ chromium (II) ion IO4¯ bromate
Mn2+ chromium (III) ion MnO4¯ hypobromite
Mn3+ manganese (II) ion IO3¯ carbonate
Cs+ manganese (III) ion OH¯ hydrogen carbonate (bicarbonate)
Cesium ion O22- sulfate
BrO3¯ hydrogen sulfate (bisulfate)
BrO¯ sulfite
CO32- hydrogen sulfite (bisulfite)
HCO3¯ oxalate
SO42- hydrogen oxalate (binoxalate)
HSO4- phosphate
SO32- hydrogen phosphate
HSO3¯ thiosulfate
C2O42- hydrogen sulfide
HC2O4-
PO43-
HPO42-
S2O32-
HS¯
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HANDOUT 2
THE VARIOUS SERIES IN ATOMIC HYDROGEN EMISSION SPECTRUM
The most intense line. (More energy is
needed to move the electron to the
higher energy level). Refer to the first
line of any series.
Converging limit or continuum
limit when n=∞
Series Region n1 n2 Spectrum
Lyman 1 2, 3, 4 ……. Ultraviolet
Balmer 2 3, 4, 5 …… Visible
Paschen 3 4, 5, 6 ……. Infrared
Brackett 4 5, 6, 7 …… Infrared
Pfund 5 6, 7, 8 …… Infrared
MARDIANA AMIR MIZAN 39
CHEMISTRY (SK015)
HANDOUT 3
THE ELECTRONEGATIVITY OF ELEMENTS AND THE GROUND STATE
ELECTRON CONFIGURATIONS OF ELEMENTS
MARDIANA AMIR MIZAN 40
CHEMISTRY (SK015)
HANDOUT 4
VALENCE SHELL ELECTRON-PAIRED REPULSION (VSEPR) THEORY
MOLECULES IN WHICH CENTRAL ATOM WITHOUT LONE PAIR
No. of atoms No. of lone Arrangement of Molecular
electron pairs geometry
Class bonded to pairs on Examples
central atom central atom
AB2 2 0 linear linear BeCl2
HgCl2
0 BF3
AB3 3 trigonal trigonal
planar planar
AB4 4 0 CH4
AB5 5
tetrahedral tetrahedral
PCl5
0
Trigonal Trigonal
bipyramidal bipyramidal
XeF6
AB6 6 0
octahedral octahedral
MARDIANA AMIR MIZAN 41
CHEMISTRY (SK015)
MOLECULES IN WHICH CENTRAL ATOM WITH LONE PAIR
No. of atoms No. of lone Arrangement of Molecular
electron pairs geometry
Class bonded to pairs on Examples
central atom central atom
AB3 3 0 BF3
1 SO2
AB2E 2 B e nt
A
<120o
BB
AB4 4 0 CH4
Trigonal pyramidal
AB3E 3 1 A B NH3
B
B
107o Trigonal pyramidal
104o
AB2E2 2 2 B e nt H2O
A
BB
AB5 5 0 PCl5
MARDIANA AMIR MIZAN 42
CHEMISTRY (SK015)
Distorted tetrahedron
@ see saw
AB4E 4 1 SF4
B AB
<90o,<120o BB
AB3E2 3 T -s ha pe d
B
2 ClF3
BA
<90o B
L inear
B
AB2E3 2 3 A I3
180o B
AB6 6 0 XeF6
Square pyramidal
BB
AB5E 5 1 A BrF5
BB XeF4
B
<90o 43
Square planar
BB
AB4E2 4 2 A
90o B B
MARDIANA AMIR MIZAN
CHEMISTRY (SK015)
HANDOUT 5
COMMON ACID-BASE INDICATORS AND RELATIVE STRENGTHS OF ACID-BASE
CONJUGATE PAIRS
1) Some Common Acid-base Indicators
Indicator Colour pH Range *
Alizarin yellow R In Acid In Base 10.1-12.0
Bromocresol green 3.8-5.4
Bromophenol blue Yellow Orange-red 3.0-4.6
Bromothymol blue 6.0-7.6
Chlorophenol blue Yellow Blue 4.8-6.4
Congo red 3.0-5.2
Cresol red Yellow Violet 7.2-8.8
Indigo carmine 11.4-13.0
Litmus (Azolitmin) Yellow Blue 4.5-8.3
Malachite green 0.2-1.8
Methyl orange Yellow Red 3.1-4.4
Methyl orange in xylene cyanole solution 3.2-4.2
Methyl red Blue Red 4.2-6.3
Methyl violet 0.0-1.6
Methyl yellow (in ethanol) Yellow Red 2.9-4.0
Neutral red 6.8-8.0
Phenolphtalein Blue Yellow 8.2-10.0
Phenol red 6.6-8.0
Thymol blue (acid) Red Blue 1.2-2.8
Thmol blue (base) 8.0-9.6
Thymolphthalein Yellow Blue-green 9.4-10.6
Red Yellow
Purple Green
Red Yellow
Yellow Blue-violet
Red Yellow
Red Yellow
Colourless Pink
Yellow Red
Red Yellow
Yellow Blue
Colourless Blue
2) Relative Strengths of Conjugate Acid-Base Pairs
Strong Acid Conjugate Base Strength
Strength acids HClO4 (perchloric acid) ClO4- (perchlorate ion) of base
of acid HBr (hydrobromic acid) increases
increases Weak HCl (hydrochloric acid) Br- (bromide ion)
H2SO4 (sulphuric acid) 44
acids Cl- (chloride ion)
HNO3 (nitric acid) HSO4- (hydrogen sulphate ion)
H3O+ (hydronium ion)
HF (hydrofluoric acid) NO3- (nitrate ion)
HCOOH (formic acid)
CH3COOH (acetic acid) H2O (water)
HCN (hydrocyanic acid) F- (fluoride ion)
HCOO- (formate ion)
NH3 (ammonia) CH3COO- (acetate ion)
CN- (cyanide ion)
NH2- (amide ion)
MARDIANA AMIR MIZAN
CHEMISTRY (SK015)
HANDOUT 6
THE ACID-BASE TITRATIONS CURVE
MARDIANA AMIR MIZAN 45
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