CAPE Chemistry Syllabus UNIT 1
page 1 of 9
Atomic Structure & Periodic Table (Module 1)
a) discuss process of theoretical change with respect to Dalton’s atomic
theory
b) know the structure of an atom with respect to (wrt) properties of the
sub-atomic particles, their relative charges and masses, location and
behaviour in electric and magnetic fields
c) define proton number, mass number, isotopes, relative atomic mass
d) state and explain concept of radioactivity
e) describe the different types of nuclear decay
f) write nuclear equations illustrating the different types of decay
g) understand the n/p ratio, the island of stability and its significance to
the stability of the nucleus
h) be able to predict which type of decay a radioactive isotope would
undergo to fall within the island of stability
i) state at least three uses of radioactive isotopes
j) calculate relative atomic masses of elements given isotopic masses and
abundances
k) describe how an emission spectrum of an element is obtained
l) show how emission spectra provides evidence for the existence of
discrete energy levels (refer to the Bohr model, using the hydrogen
spectrum as an example)
m) state and explain the origins of the Lyman, Balmer and Paschen series
in the hydrogen spectrum
n) be able to perform calculations using energy, wavelength and
frequency e.g. ΔE OR dE = hv
o) define the terms orbital, sub-shell and shell
p) state and describe the shapes of s and p orbitals and the relative
energies of s, p and d orbitals (pay special attention to 3d and 4s relative
energies)
q) state and apply the rules of filling s, p and d orbitals with electrons
r) state the electronic configurations of atoms and common ions of
elements with atomic number from 1 – 30
s) define the term first ionisation energy and illustrate it with an equation
t) state and explain the factors that influence first ionisation energy
(atomic radii, nuclear charge, shielding)
u) state and explain the general trend in the first ionisation energy in
periods 2 and 3 as well as the anomalies
v) show how first ionisation energy gives evidence for sub-shells (use
period 3)
w) be able to determine the group # of an element based on ionisation
energies
CAPE Chemistry Syllabus UNIT 1
page 2 of 9
Forces of Attraction (Module 1)
a) state the various forces of attraction between particles:- ionic bonds,
covalent bonds, metallic bonds, dative bonds, hydrogen bonds, Van der
Waals forces
b) state the relationship between the forces of attraction and states of
matter
c) relate physical properties( variation in m.p /b.p, solubilities) to
differences in strength of forces of attraction
d) explain the formation of i) ionic bonds ii) covalent bonds (sigma and
pi bonds) iii) metallic bonds, also include relationship between
electronegativity and polarity of bonds
e) describe the formation of dative bonds e.g. NH3 + H+, BF3 and NH3
f) explain how hydrogen bonds, Van der Waals forces (temporary dipoleinduced dipole and permanent dipole-dipole) are formed
g) state and explain VSEPR theory
h) use VESPR theory to predict shapes of, and bond angles in simple
molecules and ions  BeCl2 (linear), BF3 (trigonal planar), NH4+, CH4
(tetrahedral), NH3, CH3+, H3O+ (pyramidal), H2O (bent / non-linear)
i) explain concept of hybridisation and resonance
j) explain shapes and bond angles of methane, ethane, ethane and
benzene
k) predict shapes and bond angles of molecules similar to ethane and
ethane e.g. substituted derivates
l) describe qualitatively the lattice structure of crystalline solids and their
relation to physical properties e.g. simple molecular (I2), hydrogen
bonded (ice), giant molecular (SiO2), ionic (NaCl), metallic (Cu) and
giant atomic (diamond, graphite)
The Mole Concept
a) define mole, Avogadro’s Law and molar mass
b) perform calculations involving Avogadro’s Law and molar mass
c) construct balanced ionic equations from full equations
d) state molar volumes of gases at stp and rtp
e) define molar concentration and mass concentration
f) perform calculations involving molar conc., mass conc., and volumes
of gases
g) calculate empirical and molecular formulae given appropriate data
h) perform calculations involving titrimetric analyses
CAPE Chemistry Syllabus UNIT 1
page 3 of 9
Redox Reactions (Module 1)
a) define the term redox
b) define the terms oxidation and reduction
c) determine the oxidation state of free elements or elements within
compounds using oxidation state rules
d) determine whether a reaction is redox or non-redox
e) determine whether a substance is a reducing agent or oxidising agent
based on changes of oxidation states
f) write a full ionic equation from two half equations
Kinetic Theory (Module 1)
a) state basic assumptions of the kinetic theory with reference to an ideal
gas
b) explain the differences between real and ideal gases
c) describe the conditions necessary for a gas to approach ideal
behaviour
d) state why gases deviate from ideal gas behaviour at high pressures and
low temperatures
e) illustrate ideal gas behaviour using graphical representations
f) state Charles’ Law and Boyle’s Law and ideal gas equation and
perform calculations
g) represent Charles’ Law and Boyle’s Law graphically
h) explain i) the liquid state ii) melting iii) vaporisation
Energetics (Module 1)
a) state conditions needed for a reaction to occur
b) state that in all chemical reactions bonds are broken and bonds are
formed
c) understand that energy changes occur with the breakage and formation
of bonds
d) understand which changes require or release energy
e) define the terms endothermic and exothermic
f) list at least 3 examples of endothermic and exothermic reactions
g) illustrate endothermic and exothermic reactions using energy profile
diagrams
h) define the term bond energy
i) show the relationship between bond energy, strength of covalent bonds
and reactivity of covalent molecules (use lack of reactivity of nitrogen as
an example)
j) state and explain factors that affect bond energy
CAPE Chemistry Syllabus UNIT 1
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k) define the terms:- enthalpy of formation, enthalpy of combustion,
enthalpy of reaction, enthalpy of hydration, enthalpy of solution,
enthalpy of neutralisation, enthalpy of atomisation, ionisation energy,
electron affinity and lattice energy
l) states processes involved in formation of an ionic compound
m) understand that certain processes cannot be determined
experimentally and must be calculated using Hess’s Law
n) define Hess’s Law
o) understand that the Born-Haber cycle is an application of Hess’s Law
for ionic compounds
p) state steps involved in a Born-Haber cycle
q) use Hess’s Law to determine unknown processes in the formation of
an ionic compound given appropriate data
r) state and explain the effect of ionic charge and radius on the
magnitude of lattice energy
**************END OF MODULE 1************
Rates of reactions (Module 2)
a) states conditions needed for a reaction to occur
b) state and explain factors that affect rates of reactions (include simple
experiments that could study the effects of any factor)
c) explain necessity of knowledge of reaction rates e.g. keeping food in
fridge to last longer, manufacture of chemicals in shortest possible time
d) define and describe the terms in relation to reaction rates:- rate
constant, order of reaction, half-life, rate determining step, activation
energy, collision theory, catalysis
e) state and explain different types of catalysis (homogeneous and
heterogeneous)
f) illustrate zero, first and second order reactions graphically
g) determine the order of a reaction using supplied data
h) construct rate equations for zero, first and second order reactions using
data supplied from experiments
i) use order of reactions to deduce possible reaction mechanisms
j) interpret concentration vs time and concentration vs rate for zero and
first order reactions
k) calculate initial rates and rate constants using rate data
l) perform calculations using half-life data
m) state and explain the effect of temperature and catalysts on the rate of
a reaction (use Boltzman distribution curve and collision frequency)
CAPE Chemistry Syllabus UNIT 1
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Principles of Chemical Equilibrium (Module 2)
a) explain the term equilibrium in a chemical reaction
b) state differences between static and dynamic equilibrium
c) state 4 characteristics of dynamic equilibrium
d) define terms K, Kc and Kp (write equilibrium constant expressions)
e) perform calculations involving Kc and Kp
f) state Le Chatelier’s principle and apply it to explain effect of changes
in temperature, concentration and pressure on a system in equilibrium
g) state and explain how changes in concentration, pressure, temperature
or a catalyst would affect the value of the equilibrium constant (use
Haber and Contact processes as example)
Acid/Base Equilibrium (Module 2)
a) define the terms strong and weak for acids and bases using BronstedLowry theory
b) define the terms Ka, Kb, pH, pKa, pKb, Kw, pKw
c) perform calculations involving Ka, Kb, pH, pKa, pKb, Kw, pKw
d) sketch pH changes during various types of acid-base titrations
e) describe and explain pH changes of an indicator (use methyl orange
and phenolphthalein as examples)
f) state basis of choice of indicator for use in a titration
g) show and explain the effectiveness of different indicators in titrations
Buffers and pH (Module 2)
a) define the term buffer solution
b) explain how buffer solutions control pH
c) calculate the pH of buffer solutions from appropriate data
d) describe simple experiments on how to determine the pH of buffer
solutions
e) state the buffer systems in blood (hydrogen carbonate, phosphate and
amino acids) and the importance of presence of buffer systems in
biological systems
f) state how exercise can change the pH of blood
g) state and explain how the buffer systems in blood control pH during
various levels of exercise
h) explain importance of buffer systems in industrial processes e.g.
enzyme catalysed reactions and food processing industry
CAPE Chemistry Syllabus UNIT 1
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Solubility Product (Module 2)
a) define the term solubility product (write the Ksp expression for any
given sparingly soluble salt)
b) state and explain principles underlying solubility product
c) state the relationship between Le Chatelier’s Principle and the
common ion effect
d) describe how common ion effect can be illustrated via
experimentation
e) perform calculations involving solubility product
f) describe how to determine the solubility product of a substance
g) relate solubility product principle to qualitative analysis and kidney
stone formation
Redox Equilibrium (Module 2)
a) define the terms standard electrode potential and standard cell
potential
b) describe, sketch and label a diagram of the standard hydrogen
electrode
c) describe methods used to measure the standard electrode potentials
of:- metals or non-metals in contact with their ions in aqueous solution
and ions of the same element in different oxidation states (include
labelled diagrams of electrochemical cells)
d) be able to calculate standard cell potentials from standard electrode
potentials of two half cells
e) use standard electrode potentials of cells to determine:- the direction
of electron flow and the feasibility of a reaction (include cell diagram or
notation of the type Zn (s)| Zn2+ (aq) || Cu2+ (aq) | Cu (s)
f) predict how the value of an electrode potential varies with
concentration (apply Le Chatelier’s Principle)
g) apply the principles of redox processes to energy storage devices
(include references to batteries and fuel cells)
**************END OF MODULE 2************
CAPE Chemistry Syllabus UNIT 1
page 7 of 9
Period Sodium to Argon (Module 3)
a) explain the variation in physical properties of the elements in terms of
structure and bonding (include reference to melting point, electrical
conductivity, atomic radii, electronegativity and density
b) describe the reactions of the elements with oxygen, chlorine and water
(no treatment of peroxides and superoxides required)
c) explain the variation in oxidation number of the oxides and chlorides
d) describe the reactions of the oxides chlorides in water
e) explain the trend in the acid/base behaviour of the oxides and
hydroxides
f) predict the types of chemical bonding present in the chlorides and
oxides (refer to differences in electronegativities and ionic radii of the
elements)
Group II elements (Module 3)
a) explain the variations in properties of the elements in terms of
structure and bonding (include reference to atomic and ionic radii and
ionisation energies)
b) describe the reactions of the elements with oxygen, water and dilute
acids
c) explain the variation in the solubility of the sulphates (refer to lattice
and hydration energies)
d) explain the variation in the thermal decomposition of the carbonates
and nitrates
e) discuss the uses of some of the compounds of magnesium and calcium
( limited to the use of magnesium oxide, calcium oxide, calcium
hydroxide and calcium carbonate)
Group IV elements (Module 3)
a) explain the variation in physical properties of the elements in terms of
structure and bonding ( refer to variations in metallic character and
electrical conductivity)
b) describe the bonding of the tetrachlorides
c) explain the reactions of the tetrachlorides with water
d) discuss the trends in:- (i) bonding (ii) acid/base character (iii) thermal
stability of the oxides of oxidation states II and IV (make reference to Eө
values of the elements)
e) discuss the uses of ceramics based on silicon(IV) oxide
CAPE Chemistry Syllabus UNIT 1
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Group VII elements (Module 3)
a) explain the variations in physical properties of the elements in terms of
structure and bonding (refer to volatility, density, state, only a
description of the colour is required)
b) explain the relative reactivities of the elements as oxidising agents
(include reactions with sodium thiosulphate and refer to Eө values, also
use solutions of the elements with bleach, bromine water and iodine
solution)
c) describe the reactions of the elements with hydrogen
d) explain the relative stabilities of the hydrides (include reference to
bond energies)
e) describe the reactions of the halide ions with:- (i) aqueous solution of
AgNO3 followed by aqueous ammonia (ii) conc. sulphuric acid
f) describe the reactions of chlorine with cold and with hot aqueous
solution of sodium hydroxide (mention changes in oxidation number)
First Row Transition Metals Module 3)
a) describe the characteristics of transition elements ( include variation in
oxidation number, complex formation, coloured compounds, catalytic
activity, magnetic properties)
b) determine the electronic configuration of the first row of transition
elements and of their ions (include mention of changes in oxidation
number)
c) explain the relatively small changes in atomic radii, ionic radii and
ionisation energies of the elements across the period
d) explain the formation of coloured ions by transition elements (refer to
d orbital separation of energy in octahedral complexes)
e) describe the variation in oxidation states of vanadium (include the use
of an acidified solution of ammonium vanadate(V) and granulated zinc)
f) discuss qualitatively the properties of a transition element when
compared to those of calcium as a typical s-block element (refer to
melting point, density, atomic radius, ionic radius, first ionisation energy
and conductivity)
g) predict the shapes of complexes of transition elements (refer to
octahedral, tetrahedral and square planar)
h) discuss the use of Fe2+ (aq)/Fe3+ (aq), MnO4- (aq)/ Mn2+ (aq), Cr2O72(aq)/ Cr3+ (aq) as redox systems
i) explain the principle of ligand exchange (refer to stability constants
and the CO/O2 haemoglobin and NH3 (aq)/Cu2+ (aq) systems)
j) describe the process of ligand exchange with reference to Co2+ (aq) and
Cu2+ (aq)
CAPE Chemistry Syllabus UNIT 1
page 9 of 9
Identification of Cations and Anions (Module 3)
a) identify the cations K+, Na+, Ca2+, Ba2+, Cu2+ by their flame tests (refer
to atomic emission spectra in Module 1
b) identify aqueous cations Mg2+, Al3+, Ca2+, Cr3+, Mn2+, Fe2+, Fe3+, Cu2+,
Zn2+, Ba2+, Pb2+, NH4+ (include reactions with OH- (aq), CO32- (aq) and
NH3 (aq) and confirmatory tests)
c) explain the principles upon which the reactions are based (refer to
basic oxides, amphoteric oxides and complexation)
d) write ionic equations for the reactions of the cations
e) identify anions CO32-, NO3-, SO42-, SO32-, Cl-, Br-, I-, CrO42- (include
the reactions with HCl, conc. H2SO4, Pb2+ (aq), Ag+ (aq) followed by
NH3 (aq), Ca(OH)2 (aq), Ba2+ (aq) followed by dilute acid. For NO3- use
copper turnings and conc H2SO4 or add aluminium powder or zinc
powder in the alkaline solution and confirmatory tests for gases where
applicable
f) write ionic equations with state symbols for the reactions of then
anions
*********************END OF UNIT 1*********************