Chemistry - 3.1.8: Thermodynamics

Standard enthalpy of atomisation

The enthalpy change which accompanies the formation of one mole of gaseous atoms from the element in its standard state under standard conditions.

First ionisation energy

The standard enthalpy change when one mole of gaseous atoms is converted into a mole of gaseous ions, each with a single positive charge.

Symbol for standard enthalpy of atomisation

ΔH at

Symbol for first ionisation energy

IE

Symbol for electron affinity

ΔH ea

First electron affinity

The standard enthalpy change when a mole of gaseous atoms is converted to a mole of gaseous ions, each with a single negative charge.

Second electron affinity

The enthalpy change when a mole of electrons is added to a mole of gaseous ions with a single negative charge to form ions with two negative charges.

Symbol for lattice enthalpy of formation

ΔH lat

Lattice enthalpy of formation

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

When a lattice forms, what is the enthalpy change always and why?

It is always negative as new bonds are forming so energy is given out

Lattice enthalpy of dissociation

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

When a lattice dissociates, what is the enthalpy change always and why?

It is always positive as bonds are breaking so energy is being taken in

Enthalpy of hydration and symbol

The standard enthalpy change when one mole of gaseous ions are converted into aqueous ions/solution (ΔH hyd)

Enthalpy of solution and symbol

The standard enthalpy change when one mole of solute dissolves completely in sufficient solvent to infinite dilution, forming a solution in which the molecules or ions are far enough apart to not interact with each other (ΔH sol)

Are first electron affinities always positive or negative and why?

Negative - there is an attractive force between the nucleus of an atom and an external electron, so energy is given out

Are second electron affinities always positive or negative and why?

Positive - energy must be put in to overcome the repulsion between an electron and a negatively charged ion

Give the equation for the lattice enthalpy of formation of NaCl

Na+ (g) + Cl- (g) -> NaCl (s)

List the 5 energy changes that occur when measuring the enthalpy change of formation of a giant ionic lattice

1. Atomisation of metal

2. Atomisation of non-metal

3. Ionisation energies of metal

4. Electron affinities of non-metal

5. Formation of lattice

How do you draw a Born-Haber cycle? (6)

1. All elements start in standard states

2. Atomisation of metal (positive so uphill)

3. Atomisation of non-metal (positive so uphill)

4. Ionisation energies of metal (positive so uphill)

5. Electron affinity of non-metal (negative so downhill)

6. Lattice formation enthalpy (downhill as changing from separate ions to lattice)

Born-Haber cycle

A thermochemical cycle that includes all the enthalpy changes involved in the formation of an ionic compound

Do larger ions have smaller or larger lattice enthalpies and why?

Smaller - the opposite charges do not approach each other as closely when the ions are larger

Do ions with a greater charge but similar size have smaller or larger lattice enthalpies and why?

Larger - ions with a greater charge give out more energy when they come together

Hydration

A reaction in which water is added

What is the effect of charge and size on the enthalpy of hydration and why?

It becomes more negative for highly charged ions and less negative for bigger ions. The greater the charge / the smaller the size of the ion, the stronger its attraction to H / O

What 3 processes occur when dissolving an ionic compound in water?

1. Lattice dissociation enthalpy: the ionic lattice is broken to give separate gaseous ions

2. Enthalpy of hydration is given out: the positive (metal) ions are hydrated

3. Enthalpy of hydration is given out again: the negative (non-metal) ions are hydrated

Why is there sometimes a large discrepancy between the experimental (via B-H cycle) and calculated values for lattice formation enthalpy?

The bond has some covalent character (e.g. a greater experimental value means more bonding is present) due to one of the ions becoming polarised, so has stronger/weaker forces holding the lattice together.

Polarised

When the distribution of charge around an atom or ion is distorted from the spherical

What factors affect polarisation for:

1. positive ions

2. negative ions?

1. Positive ions: small size, high charge

2. Negative ions: large size, high charge

What terms are used to describe reactions which could take place of their own accord?

Feasible or spontaneous

Entropy

A numerical measure of disorder in a chemical system

Symbol for entropy and entropy change

S, ΔS

What does it mean if the values for entropy change are:

1. positive

2. negative?

1. Positive: products are more disordered than the reactants

2. Negative: reactants are more disordered than the products

Order the states of matter in ascending entropy values.

Solids < Liquids < Gases

How do you calculate entropy change?

Total entropy of products - total entropy of reactants

What two factors affect the feasibility of a chemical reaction?

The enthalpy change and the entropy change

Gibbs free energy and symbol

A measure of the amount of energy available to do useful work in a chemical reaction, taking into account the enthalpy change and entropy change (G)

What does it mean if ΔG is positive?

The reaction is not feasible

Formula for ΔG

ΔG = ΔH - TΔS (G = Gibbs free energy, H = enthalpy change in kJmol-1, T = temperature in K, S = entropy change in JK-1mol-1)

Why is it important that ΔG depends on temperature?

It means that some reactions may be feasible at one temperature and not at another

What does it mean if ΔG = 0?

The reaction is just feasible

Why does the theoretical enthalpy of lattice dissociation/formation differ from the experimental value calculated using a Born-Haber cycle? (2)

1. The ionic lattice (insert name) has covalent character - experimental lattice enthalpy value includes covalent interaction

2. The ionic lattice is not perfectly ionic - theoretical lattice enthalpy value assumes only ionic interaction

What is the value of the entropy when the temperature is 0K and why?

S = 0 as the particles are not moving, so there is no disorder (perfect order instead)

Why does entropy increase as temperature increases? (2)

• As temperature increases, the particles start to move

• This causes the disorder to increase

Why is there a bigger ΔS during boiling than during melting?

During boiling there is a bigger increase in disorder

How does the Gibbs free energy equation link to the equation for a straight line?

Straight line equation => y = mx + c Gibbs FE equation => ΔG = -ΔS(T) + ΔH ΔH = c, -ΔS = m (both ΔH and ΔS are constants)

As temperature increases, why does ΔG decrease?

ΔS gets bigger (particles move more so are more disordered)

In a graph of ΔG against T, why would the gradient change?

One of the products changes state, so the entropy data would be different

In the following reaction, why does the entropy change decrease? Mg (s) +1/2 O2 (g) -> MgO (s)

Decrease in the number of moles of gas so the system becomes more ordered

When is a reaction always feasible? (referring to ΔH and ΔS)

When ΔH is negative and ΔS is positive

What does it mean if the theoretical and experimental enthalpy of lattice formation are close in value?

The compound is (almost) perfectly ionic

What causes lattice enthalpy to increase more - size of the ion or charge of the ion?

Size of the ion

Which would have a lower enthalpy of hydration and why? - Ca2+ vs Mg2+

Ca2+ would have a lower enthalpy of hydration as it is bigger, so its more weakly attracted to the Oδ in water

MgCl2(s) + 4H2O(l) ⟶ MgCl2.4H2O(s)

When drawing the Hess diagram for the above reaction, if you have the enthalpies of solution for MgCl2 and MgCl2.4H2O, in what direction do you draw the arrows?

Downwards, to the elements