First Term

1 Atoms, Molecules and Stoichiometry

1.1 Fundamental particles of an atom

Candidates should be able to:

(a) describe the properties of protons, neutrons and electrons in terms of their relative charges and relative masses;
(b) predict the behaviour of beams of protons, neutrons and electrons in both electric and magnetic fields;
(c) describe the distribution of mass and charges within an atom;
(d) determine the number of protons, neutrons and electrons present in both neutral and charged species of a given proton number and nucleon number;
(e) describe the contribution of protons and neutrons to atomic nuclei in terms of proton number and nucleon number;
(f) distinguish isotopes based on the number of neutrons present, and state examples of both stable and unstable isotopes.

1.2 Relative atomic, isotopic, molecular and formula masses

Candidates should be able to:

(a) define the terms relative atomic mass, Ar, relative isotopic mass, relative molecular mass, Mr, and relative formula mass based on 12C;
(b) interpret mass spectra in terms of relative abundance of isotopes and molecular fragments;
(c) calculate relative atomic mass of an element from the relative abundance of its isotopes or its mass spectrum.

1.3 The mole and the Avogadro constant

Candidates should be able to:

(a) define mole in terms of the Avogadro constant;
(b) calculate the number of moles of reactants, volumes of gases, volumes of solutions and concentrations of solutions;
(c) deduce stoichiometric relationships from the calculations above.

2 Electronic Structures of Atoms

2.1 Electronic energy levels of atomic hydrogen

Candidates should be able to:

(a) explain the formation of the emission line spectrum of atomic hydrogen in the Lyman and Balmer series using Bohr’s Atomic Model.

2.2 Atomic orbitals: s, p and d

Candidates should be able to:

(a) deduce the number and relative energies of the s, p and d orbitals for the principal quantum numbers 1, 2 and 3, including the 4s orbitals;
(b) describe the shape of the s and p orbitals.

2.3 Electronic configuration

Candidates should be able to:

(a) predict the electronic configuration of atoms and ions given the proton number (and charge);
(b) define and apply Aufbau principle, Hund’s rule and Pauli Exclusion Principle.

2.4 Classification of elements into s, p, d and f blocks in the Periodic Table

Candidates should be able to:

(a) identify the position of the elements in the Periodic Table as
(i) block s, with valence shell configurations s1 and s2,
(ii) block p, with valence shell configurations from s2p1 to s2p6,
(iii) block d, with valence shell configurations from d1s2 to d10s2;
(b) identify the position of elements in block f of the Periodic Table.

3 Chemical Bonding

3.1 Ionic bonding

Candidates should be able to:

(a) describe ionic (electrovalent) bonding as exemplified by NaCl and MgCl2.

3.2 Covalent bonding

Candidates should be able to:

(a) draw the Lewis structure of covalent molecules (octet rule as exemplified by NH3, CCl4, H2O, CO2, N2O4 and exception to the octet rule as exemplified by BF3, NO, NO2, PCl5, SF6);
(b) draw the Lewis structure of ions as exemplified by SO42, CO32, NO3 and CN;
(c) explain the concept of overlapping and hybridisation of the s and p orbitals as exemplified by BeCl2, BF3, CH4, N2, HCN, NH3 and H2O molecules;
(d) predict and explain the shapes of and bond angles in molecules and ions using the principle of valence shell electron pair repulsion, e.g. linear, trigonal planar, tetrahedral, trigonal bipyramid, octahedral, V-shaped, T-shaped, seesaw and pyramidal;
(e) explain the existence of polar and non-polar bonds (including CC1, CN, CO, CMg) resulting in polar or/ and non-polar molecules;
(f) relate bond lengths and bond strengths with respect to single, double and triple bonds;
(g) explain the inertness of nitrogen molecule in terms of its strong triple bond and non-polarity;
(h) describe typical properties associated with ionic and covalent bonding in terms of bond strength, melting point and electrical conductivity;
(i) explain the existence of covalent character in ionic compounds such as A12O3, A1I3 and LiI;
(j) explain the existence of coordinate (dative covalent) bonding as exemplified by H3O+, NH4+, A12C16 and [Fe(CN)6]3.

3.3 Metallic bonding

Candidates should be able to:

(a) explain metallic bonding in terms of electron sea model.

3.4 Intermolecular forces: van der Waals forces and hydrogen bonding

Candidates should be able to:

(a) describe hydrogen bonding and van der Waals forces (permanent, temporary and induced dipole);
(b) deduce the effect of van der Waals forces between molecules on the physical properties of substances;
(c) deduce the effect of hydrogen bonding (intermolecular and intramolecular) on the physical properties of substances.

4 States of Matter

4.1 Gases

Candidates should be able to:

(a) explain the pressure and behaviour of ideal gas using the kinetic theory;
(b) explain qualitatively, in terms of molecular size and intermolecular forces, the conditions necessary for a gas approaching the ideal behaviour;
(c) define Boyle’s law, Charles’ law and Avogadro’s law;
(d) apply the pV nRT equation in calculations, including the determination of the relative molecular mass, Mr;
(e) define Dalton’s law, and use it to calculate the partial pressure of a gas and its composition;
(f) explain the limitation of ideality at very high pressures and very low temperatures.

4.2 Liquids

Candidates should be able to:

(a) describe the kinetic concept of the liquid state;
(b) describe the melting of solid to liquid, vaporisation and vapour pressure using simple kinetic theory;
(c) define the boiling point and freezing point of liquids.

4.3 Solids

Candidates should be able to:

(a) describe qualitatively the lattice structure of a crystalline solid which is:
(i) ionic, as in sodium chloride,
(ii) simple molecular, as in iodine,
(iii) giant molecular, as in graphite, diamond and silicon(IV) oxide,
(iv) metallic, as in copper;

(b) describe the allotropes of carbon (graphite, diamond and fullerenes), and their uses.

4.4 Phase diagrams

Candidates should be able to:

(a) sketch the phase diagram for water and carbon dioxide, and explain the anomalous behaviour of water;
(b) explain phase diagrams as graphical plots of experimentally determined results;
(c) interpret phase diagrams as curves describing the conditions of equilibrium between phases and as regions representing single phases;
(d) predict how a phase may change with changes in temperature and pressure;
(e) discuss vaporisation, boiling, sublimation, freezing, melting, triple and critical points of H2O and CO2;
(f) explain qualitatively the effect of a non-volatile solute on the vapour pressure of a solvent, and hence, on its melting point and boiling point (colligative properties);
(g) state the uses of dry ice.

5 Reaction Kinetics

5.1 Rate of reaction

Candidates should be able to:

(a) define rate of reaction, rate equation, order of reaction, rate constant, half-life of a first-order reaction, rate determining step, activation energy and catalyst;
(b) explain qualitatively, in terms of collision theory, the effects of concentration and temperature on the rate of a reaction.

5.2 Rate law

Candidates should be able to:

(a) calculate the rate constant from initial rates;
(b) predict an initial rate from rate equations and experimental data;
(c) use titrimetric method to study the rate of a given reaction.

5.3 The effect of temperature on reaction kinetics

Candidates should be able to:

(a) explain the relationship between the rate constants with the activation energy and temperature using Arrhenius equation k = Ae ^[-(Ea/RT)]
(b) use the Boltzmann distribution curve to explain the distribution of molecular energy.

5.4 The role of catalysts in reactions

Candidates should be able to:

(a) explain the effect of catalysts on the rate of a reaction;
(b) explain how a reaction, in the presence of a catalyst, follows an alternative path with a lower activation energy;
(c) explain the role of atmospheric oxides of nitrogen as catalysts in the oxidation of atmospheric sulphur dioxide; (d) explain the role of vanadium (V) oxide as a catalyst in the Contact process;
(e) describe enzymes as biological catalysts.

5.5 Order of reactions and rate constants

Candidates should be able to:

(a) deduce the order of a reaction (zero-, first- and second-) and the rate constant by the initial rates method and graphical methods;
(b) verify that a suggested reaction mechanism is consistent with the observed kinetics;
(c) use the half-life (t½) of a first-order reaction in calculations.

6 Equilibria

6.1 Chemical equilibria

Candidates should be able to:

(a) describe a reversible reaction and dynamic equilibrium in terms of forward and backward reactions;
(b) state mass action law from stoichiometric equation;
(c) deduce expressions for equilibrium constants in terms of concentrations, Kc, and partial pressures, Kp, for homogeneous and heterogeneous systems;
(d) calculate the values of the equilibrium constants in terms of concentrations or partial pressures from given data;
(e) calculate the quantities present at equilibrium from given data;
(f) apply the concept of dynamic chemical equilibrium to explain how the concentration of stratospheric ozone is affected by the photodissociation of NO2, O2 and O3 to form reactive oxygen radicals;
(g) state the Le Chatelier‟s principle and use it to discuss the effect of catalysts, changes in concentration, pressure or temperature on a system at equilibrium in the following examples:
(i) the synthesis of hydrogen iodide,
(ii) the dissociation of dinitrogen tetroxide,
(iii) the hydrolysis of simple esters,
(iv) the Contact process,
(v) the Haber process,
(vi) the Ostwald process;

(h) explain the effect of temperature on equilibrium constant from the equation, ln K= -H/RT + C

6.2 Ionic equilibria

Candidates should be able to:

(a) use Arrhenius, BrØnsted-Lowry and Lewis theories to explain acids and bases;
(b) identify conjugate acids and bases;
(c) explain qualitatively the different properties of strong and weak electrolytes;
(d) explain and calculate the terms pH, pOH, Ka, pKa, Kb, pKb, Kw and pKw from given data;
(e) explain changes in pH during acid-base titrations;
(f) explain the choice of suitable indicators for acid-base titrations;
(g) define buffer solutions;
(h) calculate the pH of buffer solutions from given data;
(i) explain the use of buffer solutions and their importance in biological systems such as the role of H2CO3 / HCO3 in controlling pH in blood.

6.3 Solubility equilibria

Candidates should be able to:

(a) define solubility product, Ksp;
(b) calculate Ksp from given concentrations and vice versa;
(c) describe the common ion effect, including buffer solutions;
(d) predict the possibility of precipitation from solutions of known concentrations;
(e) apply the concept of solubility equilibria to describe industrial procedure for water softening.

6.4 Phase equilibria

Candidates should be able to:

(a) state and apply Raoult’s law for two miscible liquids;
(b) interpret the boiling point-composition curves for mixtures of two miscible liquids in terms of ideal‟ behaviour or positive or negative deviations from Raoult’s law;
(c) explain the principles involved in fractional distillation of ideal and non ideal liquid mixtures;
(d) explain the term azeotropic mixture;
(e) explain the limitations on the separation of two components forming an azeotropic mixture;
(f) explain qualitatively the advantages and disadvantages of fractional distillation under reduced pressure.

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