Thermodynamics Flashcards
Enthalpy of formation
Enthalpy change when one mole of a substance is formed from its constituent elements with all substances in their standards states
Exothermic (usually)
Enthalpy of combustion
Enthalpy change when one mole of a substance undergoes complete combustion in oxygen with all substances in their standard states
Exothermic
Enthalpy of neutralisation
Enthalpy change when one mole of water is formed in a reaction between an acid and alkali under standard conditions
Exothermic
First ionisation enthalpy
Enthalpy change when each atom in a one mole of gaseous atoms loses one electron to form one mole of gaseous 1+ ions
Endothermic
First electron affinity
Enthalpy change when each atom in one mole of gaseous atoms gains one electron to form one mole of gaseous 1- ions
Exothermic for most non-metals
Second electron affinity
Enthalpy change when each atom in one mole of gaseous 1- atoms gains one electron to form one mole of gaseous 2- ions
Endothermic
Enthalpy of atomisation
Enthalpy change when one mole of gaseous atoms is produced from an elements in its standard state
Endothermic
Hydration enthalpy
Enthalpy change when one mole of gaseous ions become hydrated (converted to aqueous ions)
Exothermic
Enthalpy of solution
Enthalpy change when one mole of an ionic solid dissolves in an amount of water large enough so that the dissolved ions are well seperated and do not interact with each other
Bond dissociation enthalpy
Enthalpy change when one mole of covalent bonds is broken in the gaseous state
Endothermic
Lattice enthalpy of formation
Enthalpy change when one mole of a solid ionic compound is formed from its constituent ions in the gas phase
Exothermic
Lattice enthalpy of dissociation
Enthalpy change when one mole of a solid ionic compound is broken up into its constituent ions in the gas phase
Endothermic
Enthalpy of vapourisation
Enthalpy change when one mole of a liquid is turned into a gas
Endothermic
Enthalpy of fusion
Enthalpy change when one mole of a solid is turned into a liquid
Endothermic
What does a gap between experimental and theoretical lattice enthalpies mean?
The larger the gap, the greater the covalent character of the compound and the more distorted the ions
Are positive lattice enthalpies those of formation or dissociation?
Dissociation
Energy is taken in when electrostatic attractions between ions are broken
Are negative lattice enthalpies those of formation or dissociation?
Formation
Energy is released when electrostatic attractions between ions are formed
How are theoretical lattice enthalpies calculated?
Using the charge and size of ions and assuming the structure is perfectly ionic
Entropy calculation
Entropy change = (Sum of products entropy) - (sum of reactants entropy)
Units = Jul mol-1 K-1
Gibbs free energy calculation
Gibbs free energy = enthalpy change - (temp x entropy change)
Outline how Gibbs free energy can indicate whether or not a reaction is feasible
If G is less than or equal to zero = feasible
If G is more than zero = not feasible
What causes covalent character within ionic bonds?
- A small, highly positive ion and a large, highly negative ion
- This distorts the shape of the electron cloud in the negative ion, pulling electrons towards the positive ion
Give the equation for enthalpy of solution
Lattice dissociation enthalpy + enthalpy of hydration
Define enthalpy change
Heat change at constant pressure
Why is the enthalpy of hydration for Ca2+ less exothermic than Mg2+?
- Ca2+ is larger
- So it has a weaker attraction to the partially negative O- in water
At what temp is the entropy of any element zero?
0K
Explain why the enthalpy of hydration becomes less exothermic from Li+ to K+
- Size of ion increases
- So attraction between metal ion and Oδ– of water decreases
How would you use a graph showing Gibbs free energy against temperature to calculate delta H and delta S?
Delta H = Intercept at y (Gibbs axis)
Delta S = Opposite sign gradient
