Thermodynamics and Further Inorganic Chemistry

1. Thermodynamics and Further Inorganic Chemistry

2. Contents • Thermodynamics • Periodicity • Redox Equilibria • Transition Metals • Reactions of Inorganic Compounds in Aqueous Solution

3. Thermodynamics • Enthalpy Change • Free Energy and Entropy Change

4. Enthalpy Change (ΔH) • Enthalpy of formation • is the enthalpy change that occurs when one mole of compound in its standard state if formed from its element in their standard states under standard conditions. • Ionisation enthalpy • of an element is the enthalpy change that accompanies the removal of 1 electron in each atom of one mole of gaseous atoms • Enthalpy of atomisation • of an element is the enthalpy change that occurs when 1 mole of gaseous atoms are formed from the element in its standard state. • Bond dissociation enthalpy • the enthalpy required to break and separate one mole of bond so that the gaseous atoms exert no force on each other. • Electron affinity • of an element is the enthalpy change that accompanies the addition of one electron to each atom in one mole of gaseous atoms. • Lattice enthalpy • of an ionic compound is the enthalpy change which accompanies the formation of one mole of ionic compound from its constituent gaseous ions. • You will need to be able to use the quantities in energy cycle calculations, by applying Hess’s Law. Remember it doesn’t matter which route a chemical change takes the net energy change stays the same.

5. Free-Energy Change (ΔG) andEntropy Change (ΔS) • ΔH, whilst important, is not sufficient to explain spontaneous endothermic • The concept of increasing disorder (entropy change, ΔS) accounts for this phenomenon. • E.g. the production of CO2 from hydrogencarbonates with acid is endothermic but occur without heating, because the products are significantly more disordered than the products. • The balance between entropy and enthalpy determines the feasibility of a reaction and is called the change in free energy. • ΔG = ΔH -TΔS (derivation not required). • ΔG must be negative for a reaction to proceed spontaneously

6. Periodicity • All trends listed here are for progression from left to right of the periodic table. • Reactions with water Na and Mg • Reactivity decreases • The elements progress from metals to non-metals. • Period 3 oxides • The amount of oxygen per mole of the element increases. • The oxides go from ionic to covalent in nature • The oxides go from basic to acidic • Boiling points start high because of the strong ionic forces but become higher in the middle oxides have mixed ionic covalent nature but still form giant lattices so have high boiling points. The latter few elements form simple molecules and so have relatively low boiling points. • With water the metal oxides form alkalis, aluminium and silicon oxide are insoluble, the non metal oxides form acidic solutions

7. Periodicity • Period 3 chlorides • The amount of chlorine atoms per mole of the elements increases across the period. • Sodium and magnesium chloride are ionic where as aluminium chloride is covalent as are the chlorides of the other elements. • The boiling points decrease as the compounds progress from giant ionic in structure to simple molecular. • Ionic chlorides form neutral solution form neutral solutions in water. Covalent chlorides are hydrolysed by water and form acidic solutions.

8. Redox Equilibria • Oxidation and reduction are electron transfer reactions, • Oxidation is the loss of electrons and reduction is the gain of electrons • In redox reactions the oxidation state of the elements involved changes. • When looking at redox reactions it is sometimes useful to look at half equations involved. • Na  Na+ + e- the sodium has lost an electron and has therefore been oxidised • 2H2O + 2e- H2 + 2OH- the hydrogen has lost an electron and has therefore been reduced. • These can be combined as follows: • 2Na + 2H2O  NaOH + H2 • Half equations are also useful when examining the reactions that occur during electrolysis

9. Transition Metals • Transition metal characteristics arise from an incomplete d sub-level. • These characteristics include complex formation, high density, high melting points, formation of coloured ions, variable oxidation state and catalytic activity.

10. Variable Oxidation States • Transition elements are able to form more than one ion, each with a different oxidation state, by losing the 4s electrons and different numbers of 3d electrons. • When forming ions, the 4s electrons are lost first, before the 3d electrons.

11. Transition Metal Complexes • A ligand is a molecule or ion that bonds to a metal ion • By forming a coordinate bond (one in which both bonding electrons come from the same element) • A complex ion is a central metal ion surrounded by ligands. • The coordination is the total number of coordinate bonds from ligands to the central transition metal ion.

12. Shapes of Molecules • Six coordinate complexes form octahedral shapes. • Four coordinate complexes form tetrahedral shapes

13. The effects of partially filled d-orbitals • In the partially filled d-orbitals of a transition metal complex there is an energy gap between the filled and unfilled orbitals. Visible light is of the right energy to promote an electron from one energy level to the next. • The partially filled orbitals in transition metals allow them to change oxidation states readily and thus facilitate the progress of some reactions

14. Reactions of Inorganic Compounds in Aqueous Solution • Lewis Acids and Lewis Bases • Metal Aqua ions • Substitution reactions

15. Lewis Acids and Lewis Bases • Lewis acids are lone pair acceptors such as the central metal ion in transition metal ion. • Lewis bases are lone pair donors such as ligands. • These definitions of acids and bases extend acid -base chemistry to reactions which do not involve proton exchange.

16. Metal Aqua Ions • Metal aqua ions are formed in aqueous solution: • [M(H2O)6]2+ and [M(H2O)6]3+ • These aqua ions can be present in the solid state for example copper sulphate is only blue as its aqua ion, in its anhydrous state it is white.

17. Substitution Reactions • The ligands NH3 and H2O are similar in size and are uncharged, therefore ligand exchange between these two ligands occurs without change of co-ordination number. • The Cl− ligand is larger than these uncharged ligands and that ligand exchange can involve a change of co-ordination number. • Substitution of unidentate ligands (ligands which forms one bond with the central metal) with a bidentate or a multidentate ligand leads to a more stable complex. • This is because there will be more ions in solution therefore higher entropy. • [M(U)6]2+ + 3B Â [M(U)3]2+ + 6B 4 ions in solution 7 ions in solution less entropy more entropy

18. Summary • Thermodynamics • Standard enthalpy changes can be used to calculate energy changes in unknown reactions • Periodicity • The trends in period three reactions can be explained in terms of the structure and bonding of their compounds • Redox Equilibria • When looking at redox reactions it is sometimes useful to write half equations. • Transition Metals • The chemistry and properties of transition metals can be explained by their partially filled d-orbitals. • Reactions of Inorganic Compounds in Aqueous Solution • The idea of aid base chemistry can be extended by the concept of Lewis acids and bases.