Enthalpy And Entropy

The enthalpy change when 1 mole of a product is formed from it’s reactants under standard conditions

The enthalpy change when 1 mole of gaseous atoms is formed from the element under standard conditions O(g) -> ½O2(g)

The energy required to remove 1 mole of electrons from 1 mole of gaseous atoms to form 1 mole of gaseous +1 ions

The energy required to remove 1 mole of electrons from 1 mole of gaseous 1+ ions to form 1 mole of gaseous 2+ ions

The enthalpy change when 1 mole of a crystallised ionic lattice is formed from it’s constituent gaseous ions

The energy required when 1 mole of gaseous atoms gain 1 mole of gaseous electrons to form 1 mole of gaseous -1 ions

The energy change when 1 mole of gaseous -1 ions gains an additional electron per atom to produce gaseous -2 ions

The enthalpy change when 1 mole of gaseous ions become aqueous ions

The enthalpy change when 1 mole of a solute dissolves in water

Always exothermic

• Ionic radius. As ionic radius increases, lattice enthalpy becomes less negitive as the attraction between the ions decreases. • Ionic charge. As ionic charge increases, lattice enthalpy becomes more negitive as the attraction between ions increases.

Sum of all other values found in the Born-Haber cycle

• Elements in standard states • Atomisation(s) - Converting elemets from their standard state into a gaseous state. Di-atomic gases should be turned into singlar molecules e.g. Cl2 -> 2Cl • Ionisation energy - allowing the gaseous elements to lose electrons. The element losing should not be ionised yet. It should be written as the element + e-. e.g. Rb(g) + e- + Cl(g). • Electron affinity - ionising the gaseous elements. • Lattice enthaply - forming the ionic solid

It can be both

Always exothermic

• The very top of the cycle should be the ionised substances forming the lattice, in a gaseous state. • The very bottom of the cycle should be the lattice compound • From the top line arrows should be pointing down for hydration enthalpy - turning gaseous ions into aqueous ions. • An arrow from the bottom line pointing to the final hydration line is solution enthalpy - the solute dissolving in water.

• Ionic radius. As ionic radius increases, hydration enthalpy becomes less negitive as the attraction between the ion and water decreases • Ionic charge. As ionic charge increases, hydration enthalpy becomes more negitive as the attraction between the ion and water increases.

If hydration enthalpy is greater than lattice enthalpy, dissolving is more happy to happen

J/K⁻¹ mol⁻¹

The dispersal of energy into a system, which increases as the system is more disordered

Sum of product entropies - sum of reactant entropies

ΔG = ΔH - TΔS • Temprature (T) must be in Kelvin • The ΔS value must be divided by 1000 to put it into kJ

If ΔG is negitive, a reaction is feasible

T = ΔH/ΔS