3.5 Unit 5 CHEM5 Energetics, Redox and Inorganic Chemistry

3.5.1 Thermodynamics

Enthalpy change (ΔH)

• 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 Born-Haber cycles to calculate lattice enthalpies from experimental data.
• be able to compare lattice enthalpies from Born.Haber cycles with those from calculations based on a perfect ionic model to provide evidence for covalent character in 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 ΔH for other reactions
• be able to explain why values from mean bond enthalpy calculations differ from those determined from enthalpy cycles

Free-energy change (ΔG) and entropy change (ΔS)

• understand that ΔH, whilst important, is not sufficient to explain spontaneous change (e.g. spontaneous endothermic reactions)
• understand that the concept of increasing disorder (entropy change ΔS) accounts for the above defi ciency, illustrated by physical change (e.g. melting, evaporation) and chemical change (e.g. dissolution, evolution of CO2 from hydrogencarbonates with acid)
• be able to calculate entropy changes from absolute entropy values
• understand that the balance between entropy and enthalpy determines the feasibility of a reaction; know that this is given by the relationship ΔG = ΔH . TΔS (derivation not required).
• be able to use this equation to determine how ΔG varies with temperature
• be able to use this relationship to determine the temperature at which a reaction is feasible

3.5.2 Periodicity

Study of the reactions of Period 3 elements Na . Ar to illustrate periodic trends

• be able to describe trends in the reactions of the elements with water, limited to Na and Mg
• 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.

A survey of the acid base properties of the oxides of Period 3 elements

• be able to explain the link between the physical properties of the highest oxides of the elements Na to S in terms of their structure and bonding
• be able to describe the reactions of the oxides of the elements Na to 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

3.5.3 Redox Equilibria

Redox equations

• be able to apply the electron transfer model of redox, including oxidation states and half equations to d block elements

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
• 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 1.00 mol dm.3 solution of ions

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

Electrochemical cells

• appreciate that electrochemical cells can be used as a commercial source of electrical energy
• appreciate that cells can be non-rechargeable (irreversible), rechargeable and fuel cells
• be able to use given electrode data to deduce the reactions occurring in non-rechargeable and rechargeable cells and to deduce the e.m.f. of a cell
• understand the electrode reactions of a hydrogen-oxygen fuel cell and appreciate that a fuel cell does not need to be electrically recharged
• appreciate the benefi ts and risks to society associated with the use of these cells

3.5.4 Transition Metals

General properties of transition metals

• know that transition metal characteristics of elements Ti to Cu arise from an incomplete d sub-level in atoms or ions
• know that these characteristics include complex formation, formation of coloured ions, variable oxidation state and catalytic activity

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 C2O4 ) or multidentate (e.g. EDTA4. )
• know that haem is an iron(II) complex with a multidentate ligand

Shapes 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 square planar complexes are also formed, e.g. cisplatin
• know that Ag+ commonly forms the linear complex [Ag(NH3)2]+ as used in Tollens' reagent

Formation of coloured ions

• know that transition metal ions can be identified by their colour, limited to the complexes in this unit
• 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: ΔE = hƒË
• appreciate that this absorption of visible light is used in spectrometry to determine the concentration of coloured ions

Variable oxidation states

• know that transition elements show variable oxidation states
• know that Cr3+ and Cr2+ are formed by reduction of Cr2O7 by zinc in acid solution
• know the redox titration of Fe2+ with MnO4 and Cr2O7 in acid solution
• be able to perform calculations for this titration and for others when the reductant and its oxidation product are given
• know the oxidation of Co2+ by air in ammoniacal solution
• know the oxidation in alkaline solution of Co2+ and Cr3+ by H2O2

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 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)
• understand how V2O5 acts as a catalyst in the Contact Process
• know that a Cr2O3 catalyst is used in the manufacture of methanol from carbon monoxide and hydrogen
• know that Fe is used as a catalyst in the Haber Process
• know that catalysts can become poisoned by impurities and consequently have reduced effi ciency; know that this has a cost implication (e.g. poisoning by sulfur 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 S2O8 catalysed by Fe2+ and autocatalysis by Mn2+ in reactions of C2O4 with MnO4).

Other applications of transition metal complexes

• understand the importance of variable oxidation states in catalysis; both heterogeneous and homogeneous catalysts
• 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
• appreciate the benefi ts and risks associated with this drug
• understand that [Ag(NH3)2]+ is used in Tollens' reagent to distinguish between aldehydes and ketones

3.5.5 Reactions of Inorganic Compounds in Aqueous Solution

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

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, Cr and Fe

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 (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 CO3
• 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 2CrO4 + 2H+ Cr2O7 + H2O

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 of unidentate ligand 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