AQA Chemistry - Thermodynamics

Lattice enthalpy of formation

Lattice enthalpy of formation

The standard enthalpy change when one mole of solid ionic compound is formed from its gaseous ions.

Lattice enthalpy of dissociation

The standard enthalpy change when one mole of solid ionic compound dissociated into its gaseous ions.

Standard enthalpy of formation

The enthalpy change when one mole of a compound is formed from its elements in their standard states, under standard conditions

First ionisation energy

Enthalpy change when one mole of electrons is removed from one from one mole of gaseous atoms to form one mole of gaseous 1+ ions

Second ionisation energy

Enthalpy change when one mole of electrons is removed from one mole of gaseous atoms 1+ ions to form one mole of gaseous 2+ ions

What does Hess's Law state?

The enthalpy change for a reaction is independent of the route taken.

What is the standard enthalpy of an element?

Zero by definition

Standard enthalpy change of combustion

The enthalpy change that takes place when one mole of a substance is completely burnt in (excess) oxygen

Standard enthalpy of atomisation

Enthalpy change when one mole of gaseous atoms is formed from a compound in it's standard state under standard conditions

First electron affinity

Enthalpy change when one mole of gaseous atoms gains one mole of electrons to form one mole of gaseous 1- ions

Second electron affinity

Enthalpy change when one mole of gaseous 1- ions gains one mole of electrons to form one mole of gaseous 2- ions

Enthalpy of hydration

Enthalpy change when one mole of gaseous ions become hydrated/dissolved in water to infinite dilution (water molecules totally surround the ion)

Mean Bond Dissociation Enthalpy

Enthalpy change when one mole of covalent bonds is broken, with all species in the gaseous state.

Born-Harber cycle

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

What factors affect the lattice Enthalpy of an ionic compound?

Size of the ions, charge on the ions

Enthalpy of solution

Enthalpy change when one mole of solute dissolves completely in a solvent to infinite dilution

What is the perfect ionic model?

Assumes that ions are perfectly spherical and that there is an even charge distribution (100% polar bonds). Act as point charges.

Why is the perfect ionic model inaccurate?

Polarisation often occurs when small positive ions or large negative ions are involved, so the ionic bond gains covalent character. Some lattices are not regular and the crystal structure can differ.

Which kind of bonds will be the most ionic? Why?

Between large positive ions and small negative ions.

Spontaneous and Feasible

Reaction does not take account of rate of reaction.

Is a reaction with a positive or negative enthalpy change more like to be spontaneous?

Negative - exothermic

Entropy

Disorder of a system. Higher value for entropy = more disordered

Units for entropy

JK-1mol-1

Second law of thermodynmaics

Entropy always increases, as it's overwhelmingly more likely for molecules to be disordered than ordered.

Is a reaction with positive or negative entropy change more likely to be spontaneous?

Positive - reactions always try and increase the amount of disorder

How would you calculate entropy change for a reaction?

Entropy change = sum of products' entropy - sum of reactants' entropy

How can you increase the lattice enthalpy of a compound? Why does it increase?

Smaller ions, as the charge centres will be closer together. Increased charge, since there will be a greater electrostatic force of attraction between the oppositely charged ions. Increasing the charge on the anion has the effect of increasing ionic size which has a smaller effect than increasing the charge on a cation.

Gibbs free energy using an equation

∆G = ∆H - T∆S

Units for Gibbs free energy

-change in Gibbs Free Energy = ∆G -change in enthalpy = ∆H -temperature in Kelvin = T -change in entropy = ∆S

What is the value of gibbs free energy for a reaction to show?

If G<0, reaction is feasible. If G=0, reaction is JUST feasible. If G>0, reaction is not feasible.

What is the significance of the temperature at which G=0?

This is the temperature (in kelvin) at which the reaction becomes feasible

How would you calculate the temperature at which a reaction becomes feasible?

T = ∆H/∆S Since G=0

What are the limitations of using G as an indicator of weather a reaction will occur?

Gibbs free energy only indicates if a reaction is feasible. It does not take into account the rate of reaction. Many reactions that are feasible at a certain temperature have a rate of reaction that is so slow that effectively no reaction is occuring.

If the reaction is exothermic and entropy increases, what is the value of G and what does this mean?

G is always negative, so reaction always feasible - product favoured

If the reaction is endothermic and entropy decreases, what is the value of G and what does this mean?

G is always positive, so reaction is never feasible - reactant favoured

If the reaction is exothermic and entropy decreases, what is the value of G and what does this mean?

Temperature dependent

If the reaction is endothermic and entropy increases, what is the value of G and what does this mean?

Temperature dependent

Why is entropy zero at 0K

No disorder - molecules/atoms are not moving or vibrating and cannot be arranged in any other way. Maximum possible state of order

What are the 2 key things to look out for to decide if entropy increases/decreases/stays constant?

-Number of moles - more moles made = increases in entropy -Going from solid = liquid/gas or liquid = gas

How is it possible for the temperature of a substance undergoing an endothermic reaction to stay constant?

The heat is given out escapes to the surroundings