A2 Module 5

Thermodynamics and Further

Inorganic Chemistry

Introduction

Energetics, introduced in the first of the foundation modules, is

extended into thermodynamics by the introduction of entropy and

free energy. Chemical properties of elements and compounds of

Period 3 are studied to illustrate periodic trends. The study of redox

chemistry reactions in AS2 is extended to include electrode potentials

and their use to predict the direction of simple redox reactions. The

characteristic properties of transition metal complexes are studied

including their use in industry, as catalysts and in medicine.

The reactions of metal ions in aqueous solution are systematised

through an understanding of hydrolysis and substitution reactions of

selected metal aqua ions.

Wherever possible, candidates should carry out experimental work to

illustrate the theoretical principles included in this module.

Candidates should:

14.1 Thermodynamics

14.1.1 Enthalpy change (DH) be able to define and apply the terms enthalpy of

formation, ionisation enthalpy, enthalpy of atomisation of an element and of a compound, bond dissociation enthalpy, electron affinity, lattice enthalpy (defined as either lattice dissociation or lattice formation), enthalpy of hydration and enthalpy of solution.

be able to construct a Born�Haber cycle for the formation of simple ionic compounds.

be able to calculate enthalpies of solution for ionic compounds from lattice enthalpies and enthalpies of hydration.

be able to use mean bond enthalpies to calculate an approximate value of DH for other reactions.

be able to explain why values from mean bond enthalpy calculations differ from those determined from enthalpy cycles.

 

 

 

14.1.2 Free energy change DG and understand that .H, whilst important, is not

entropy change DS sufficient to explain spontaneous change (e.g. spontaneous endothermic reactions).

understand that the concept of increasing disorder (entropy change DS) accounts for the above deficiency, illustrated by physical change (e.g. melting, evaporation) and chemical change (e.g. dissolution, evolution of CO 2 from hydrogencarbonates with acid).

understand that the balance between entropy and enthalpy determines the feasibility of a reaction; know that this is given by the relationship DGɵ = DHɵ � TDSɵ (derivation not required).

be able to calculate entropy changes from absolute entropy values.

14.2 Periodicity

14.2.1 Study of the reactions of be able to describe trends in the reactions of the

Period 3 elements Na � Ar to elements with water, limited to Na and Mg.

illustrate periodic trends be able to describe the trends in the reactions of

the elements Na, Mg, Al, Si, P and S with oxygen, limited

to the formation of Na2O, MgO, Al2O3, SiO2, P4O10 and

SO2.

be able to describe the trends in the reactions of the elements Na, Mg, Al, Si and P with chlorine, limited to the formation of NaCl, MgCl2, AlCl3, SiCl4 and PCl5.

14.2.2 A survey of the acid-base understand the link between the physical properties of the oxides of properties of the highest

Period 3 elements oxides of the elements Na � S and their structure and bonding.

be able to describe the reactions of the oxides of the elements Na�S with water, limited to Na2O, MgO, Al2O3, SiO2, P4O10, SO2 and SO3.

know the change in pH of the resulting solutions across the Period.

be able to explain the trends in these properties in terms of the type of bonding present.

be able to write equations for the reactions which occur between these oxides and given simple acids and bases.

 

 

14.2.3 A survey of the reactions of understand the link between the physical

the chlorides of Period 3 properties of the chlorides of the elements Na�P

elements with water and their structure and bonding.

be able to describe the reactions of the chlorides of the elements Na�P with water, limited to NaCl, MgCl2, AlCl3, SiCl4 and PCl5.

know the change in pH of the resulting solutions across the Period.

be able to explain the trends in these properties in terms of the type of bonding present.

14.3 Redox Equilibria

14.3.1 Variable oxidation state understand oxidation and reduction as electron

transfer reactions applied to reactions of d block elements.

know and be able to apply the rules for assigning oxidation states in order to work out the oxidation state of an element in a compound from its formula.

understand that changes in oxidation state involve redox processes.

be able to write half-equations identifying the oxidation and reduction processes in redox reactions when the reactants and products are specified.

be able to combine half-equations to give an overall redox equation.

14.3.2 Electrode potentials know the IUPAC convention for writing half-equations

for electrode reactions.

know and be able to use the conventional representation of cells.

understand how cells are used to measure electrode potentials by reference to the standard hydrogen electrode and know that secondary standards are normally used.

know the importance of the conditions when measuring the electrode potential, E (Nernst equation not required).

know that standard electrode potential, E , refers to conditions of 298 K, 100 kPa and 1 M solution of ions.

 

14.3.3 Electrochemical series know that standard electrode potentials can be listed

as an electrochemical series.

be able to use E values to predict the direction of simple redox reactions and to calculate the e.m.f of a cell.

14.4 Transition Metals

14.4.1 General properties of know that transition metal characteristics of

transition metals elements Sc � Cu arise from an incomplete d sub-

shell in atoms or ions.

know that these characteristics include complex formation, formation of coloured ions, variable oxidation state and catalytic activity.

14.4.2 Complex formation be able to define the term ligand.

know that co-ordinate bonding is involved in complex formation.

understand that a complex is a central metal ion surrounded by ligands.

know the meaning of co-ordination number.

understand that ligands can be unidentate (e.g. H2O, NH3 and Cl - ) or bidentate (e.g. NH2CH2CH2NH2 and ) or multidentate (e.g. EDTA4-).

know that haem is an iron(II) complex with a multidentate ligand.

14.4.3 Shape of complex ions know that transition metal ions commonly form

octahedral complexes with small ligands (e.g. H2O and NH3).

know that transition metal ions commonly form tetrahedral complexes with larger ligands (e.g. Cl- ).

know that Ag+ commonly forms linear complexes,

(e.g. [Ag(NH3)2]+ , [Ag(S2O3)2]3 and [Ag(CN)2]- ).

14.4.4 Formation of coloured ions know that transition metal ions can be identified

by their colour, limited to the complexes in this module.

know that colour changes arise from changes in oxidation state, co-ordination number and ligand.

know that colour arises from electronic transitions from the ground state to excited states: DE = hv.

know the use of ultraviolet and visible spectrophotometry in determining the concentration of metal ions in solution after the addition of a suitable ligand to intensify the colour.

14.4.5 Variable oxidation states know that transition elements show variable

oxidation states.

know that VO2+ , V3+ and V2+ are formed by reduction of VO2+ by zinc in acid solution.

know that Cr3+ and Cr2+ are formed by reduction of Cr2O72- by zinc in acid solution.

know the redox titrations of Fe2+ with MnO4- and Cr2O72- in acid solution.

be able to perform calculations for these titrations and for others when the reductant and its oxidation product are given.

know the oxidation of Co2+ by air in ammoniacal solution.

know the oxidations in alkaline solution of Co2+ and Cr3+ by H2O2.

14.4.6 Catalysis know that transition metals and their compounds can act

as heterogeneous and homogeneous catalysts.

Heterogeneous know that a heterogeneous catalyst is in a different phase

from the reactants and that the reaction occurs at the surface.

understand that adsorption of reactants at active sites on the surface may lead to catalytic action.

know that the strength of adsorption helps to determine the activity (e.g. W too strong adsorption, Ag too weak adsorption, and hence the utility of Ni and Pt).

understand the use of a support medium to maximise the surface area and minimise the cost (e.g. Rh on a ceramic support in catalytic converters).

know that V2O5 is used as a catalyst in the Contact Process.

know that Fe is used as a catalyst in the Haber Process.

know that catalysts can become poisoned by impurities and consequently have reduced efficiency; know that this has a cost implication (e.g. poisoning by sulphur in the Haber Process and by lead in catalytic converters in cars).

Homogeneous know that when catalysts and reactants are in the same

phase, the reaction proceeds through an intermediate species (e.g. the reaction between I- and S2O82- catalysed by Fe2+ and autocatalysis by Mn2+ in titrations of C2O42- with MnO4-).

14.4.7 Other applications of understand the importance of variable oxidation states in

transition metal complexes catalysis; both heterogeneous and homogeneous

catalysts (e.g. V2O5 in the Contact Process and autocatalysis by Mn2+ in MnO4- titrations).

understand that Fe(II) in haemoglobin enables oxygen to be transported in the blood, and why CO is toxic.

know that the Pt(II) complex cisplatin is used as an anticancer drug.

understand that [Ag(NH3)2]+ is used in Tollen�s reagent and that [Ag(S2O3)2]3- is formed in photography.

know that [Ag(CN)2]- is used in electroplating.

14.5 Reactions of Inorganic

Compounds in Aqueous

Solution

14.5.1 Lewis acids and bases know the definitions of a Lewis acid and Lewis base;

understand the importance of lone pair electrons in co-ordinate bond formation.

14.5.2 Metal-aqua ions know that metal�aqua ions are formed in aqueous solution:

[M(H2O)6]2+ , limited to M = Fe, Co and Cu ;

[M(H2O)6]3+ , limited to M = Al, V, Cr and Fe.

know that these aqua ions can be present in the solid state (e.g. FeSO4.7H2O and CoCl2.6H2O).

14.5.3 Acidity or hydrolysis reactions understand the equilibria

[M(H2O)6]2+ + H2O [M(H2O)5(OH)]+ + H3O+ and

[M(H2O)6]3+ + H2O [M(H2O)5(OH)]2+ + H3O+ to show generation of acidic solutions with M3+ , and very weakly acidic solutions with M2+ .

understand that the acidity of [M(H2O)6]3+ is greater than that of [M(H2O)6]2+ in terms of the polarising power (charge/size ratio) of the metal ion.

be able to describe and explain the simple test-tube reactions of M2+ (aq) ions, limited to M = Fe, Co and Cu, and of M3+ (aq) ions, limited to M = Al, Cr and Fe, with the bases OH- , NH3 and CO32-.

know that MCO3 is formed but that M2(CO3)3 is not formed.

know that some metal hydroxides show amphoteric character by dissolving in both acids and bases (e.g. hydroxides of Al3+ and Cr3+).

know the equilibrium reaction

2CrO42- + 2H+ Cr2O72- + H2O

14.5.4 Substitution reactions understand that the ligands NH3 and H2O are similar in

size and are uncharged, and that ligand exchange occurs without change of co-ordination number (e.g. Co2+ and Cr3+).

know that substitution may be incomplete (e.g. the formation of [Cu(NH3)4(H2O)2]2+).

understand that the Cl- ligand is larger than these uncharged ligands and that ligand exchange can involve a change of co-ordination number (e.g. Co2+ and Cu2+).

know that substitution with a bidentate or a multidentate ligand leads to a more stable complex.

understand this chelate effect in terms of a positive entropy change in these reactions.