PHYSICAL CHEM - THERMODYNAMICS

• enthalpy change when 1 mole of a compound is completely burnt in O2 with all reactants and products in standard states under standard conditions • CH4 (g) + 2O2 (g) –> CO2 (g) + H2O (l)

• enthalpy change when 1 mole of a compound is formed from its constituent elements with all reactants and products in standard states under standard conditions • eg: H2 (g) + 1/2 O2 (g) –> H2O (l)

• enthalpy change when 1 mole of gaseous atoms are formed from the element in its standard state under standard conditions • eg: 1/2 O2 (g) –> O (g)

• enthalpy change when 1 mole of e^- are removed from 1 mole of gaseous atoms to form 1 mole of gaseous ions with a 1+ charge under standard conditions • eg: Mg (g) –> Mg^+ (g) + e^-

• enthalpy change when 1 mole of e^- are removed from 1 mole of gaseous 1+ ions to form 1 mole of gaseous ions with a 2+ charge under standard conditions • eg: Mg^+ (g) –> Mg^2+ (g) + e^-

• enthalpy change when 1 mole of e^- are added to 1 mole of gaseous atoms to form 1 mole of gaseous ions with a 1- charge under standard conditions • eg: O (g) + e^- –> O^- (g)

• enthalpy change when 1 mole of e^- are added to 1 mole of gaseous 1- ions to form 1 mole of gaseous ions with a 2- charge under standard conditions • eg: O^- (g) + e^- –> O^2- (g)

• enthalpy change when 1 mole of a solid ionic compound is formed from its gaseous ions under standard conditions • eg: Na^+ (g) + Cl^- (g) –> NaCl (s)

• enthalpy change when 1 mole of solid ionic compound dissociates into its gaseous ions under standard conditions • eg: NaCl (s) –> Na^+ (g) + Cl^- (g)

• enthalpy change when water molecules surround 1 mole of gaseous ions to form aq ions under standard conditions • eg: Na^+ + aq –> Na^+ (aq)

• enthalpy change when 1 mole of solute dissolves completely in sufficient solvent to form a solution in which the molecules or ions are far enough apart to not interact with each other under standard conditions • eg: NaCl (s) + aq –> Na^+ (aq) + Cl^- (aq)

• enthalpy change required to break 1 mole of a given covalent bond under standard conditions averaged over a range of compounds • eg: H2 (g) –> 2H (g)

cycle showing all the enthalpy changes involved in the formation of an ionic compound starting with elements in their standard states (0)

theoretical lattice enthalpies use the perfect ionic model which assumes: - ions are spherical - bonding is 100% ionic (no covalent character)

• ions are not perfectly spherical • polarisation often occurs when small +ve ions/large -ve ions are involved so the ionic bond gains covalent character • the more polarisation, the more covalent the bond will be

• bonds between large +ve ions and small -ve ions • a large cation has a lower charge density and therefore less ability to distort the electron cloud of the anion • a small anion holds its electrons more tightly to its nucleus, making its electron cloud less susceptible to distortion by the cation

• size of the ions • charge on the ions

• increase the charge on the ions - greater ionic charges lead to stronger electrostatic attraction • decrease the size of the ions - smaller ionic radii allow ions to pack more closely together

• bonds between the ions break (endothermic) • the aq ions are surrounded by water molecules which have a permanent dipole due to polar O-H bond • the dipole means that the δ+ H can form bonds with the -ve ions and the δ- O can form bonds with the +ve ions • bonds between ions and water are made (exothermic)

• put ionic lattice and dissolved ions on the top and connect by enthalpy change of solution (this is the direct route) • connect ionic lattice to gaseous ions by the lattice dissociation enthalpy • connect the gaseous ions to the dissolved ions by the hydration enthalpies of each ion (this is the indirect route)

• a measure of the number of ways that particles can be arranged and the number of ways that energy can be shared between them • the more disordered the particles are, the higher the entropy is (a large +ve value of entropy shows a high level of disorder)

• physical state • no. of particles

• solids have the lowest entropy as there’s hardly any disorder • gases have the highest entropy as they have the most disordered arrangements of particles

• more particles means more entropy • this is because more particles means theres more ways that they and their energy can be arranged