Energetics

Enthalpy of formation (ΔH_f^⦵ )

Enthalpy change when one mole of a substance is formed from its constituent elements with all substances in their standard states: Exothermic (-)

e.g. Na₂O(s)

2 Na(s) + ½ O₂(g) → Na₂O(s)

Enthalpy of combustion (ΔH_c^⦵ )

Enthalpy change when one mole of a substance undergoes complete combustion in oxygen with all substances in their standard states: Exothermic (-)

e.g. hydrogen

H₂(g) + ½ O₂(g) → H₂O(l)

Enthalpy of neutralisation (ΔH_neut^⦵ )

Enthalpy change when 1 mole of water is formed in a reaction between an acid and alkali under standard conditions: Exothermic (-)

e.g. H₂SO₄ + NaOH

½ H₂SO₄(aq) + NaOH(aq) → ½ Na₂SO₄(aq) + H₂O(l)

(First) Ionisation enthalpy (ΔH_IE^⦵ )

Enthalpy change when each atom in one mole of gaseous atoms loses one electron to form one mole of gaseous 1+ ions: Endothermic (+)

e.g. magnesium

Mg(g) → Mg⁺(g) + e^-

(Second) Ionisation enthalpy (ΔH_IE^⦵ )

Enthalpy change when each ion in one mole of gaseous 1+ ions loses one electron to form one mole of gaseous 2+ ions: Endothermic (+)

e.g. magnesium

Mg⁺(g) → Mg²⁺(g) + e^-

(First) Electron affinity (ΔH_EA^⦵ )

Enthalpy change when each atom in one mole of gaseous atoms gains one electron to form one mole of gaseous 1- ions: Exothermic (-) for many non metals

e.g. oxygen

O(g) + e^- → O^-(g)

(Second) Electron affinity (ΔH_EA^⦵ )

Enthalpy change when each ion in one mole of gaseous 1- ions gains one electron to form one mole of gaseous 2- ions: Endothermic (+) as adding negative electron to negative ion

e.g. oxygen

O^- (g) + e^- → O^2-(g)

Enthalpy of atomisation (ΔH_atm^⦵ )

Enthalpy change when one mole of gaseous atoms is produced from an element in its standard state: Endothermic (+)

e.g. iodine

½ I₂(s) → I(g)

Hydration enthalpy (ΔH_hyd^⦵ )

Enthalpy change when one mole of gaseous ions become hydrated (dissolved in water): Exothermic (-)

e.g. magnesium ions

Mg²⁺(g) → Mg²⁺(aq)

Enthalpy of solution (ΔH_sol^⦵ )

Enthalpy change when one mole of an ionic solid dissolves in an amount of water large enough so that the dissolved ions are separated and do not interact with each other: Varies

e.g. magnesium chloride

MgCl₂(s) → Mg²⁺(aq) + 2 Cl^-(aq)

Bond dissociation enthalpy (ΔH_BDE^⦵ )

Enthalpy change when one mole of covalent bonds is broken in the gaseous state: Endothermic (+)

e.g. I-I bond

I₂(g) → 2 I(g)

Lattice enthalpy of formation (ΔH_LE^⦵ )

Enthalpy change when one mole of a solid ionic compound is formed from its constituent gaseous ions: Exothermic (-)

e.g. magnesium chloride

Mg²⁺(g) + 2 Cl^- (g) → MgCl₂(s)

Lattice enthalpy of dissociation (ΔH_LE^⦵ )

Enthalpy change when one mole of a solid ionic compound is broken up into its constituent gaseous ions: Endothermic (+)

e.g. magnesium chloride

MgCl₂(s) → Mg²⁺(g) + 2 Cl^- (g)

Enthalpy of vaporisation (ΔH_vap^⦵ )

Enthalpy change when one mole of a liquid is turned into a gas: Endothermic (+)

e.g. H₂O(l)

H₂O(l) → H₂O(g)

Enthalpy of fusion (ΔH_fus^⦵ )

Enthalpy change when one mole of a solid is turned into a liquid: Endothermic (+)

e.g. Mg(s)

Mg(s) → Mg(l)

Mean bond enthalpy

• Energy needed to break one mole of covalent bonds in the gaseous state

• Averaged over a range of different compounds/molecules

State the assumptions made during calculations in calorimetry. (3 marks)

• All heat released from the reaction was absorbed by the solution

• Reaction goes to completion

• Specific heat capacity of solution is 4.18 J^-1 g^-1

The enthalpy of combustion determined experimentally is less exothermic than that calculated using enthalpies of formation.

Give four possible reasons for this. (4 marks)

• Heat escaped to surroundings

• Incomplete combustion of (substance)

• Evaporation of (substance) between weighing and setting light to it

• Didn’t stir to make temperature the same throughout

Define the term enthalpy change

Heat energy change at a constant pressure

Suggest how the students’ method, and the analysis of the results, could be improved in order to determine a more accurate value for the enthalpy of reaction. Justify your suggestions.

Do not refer to the precision of the measuring equipment. Do not change the amounts or the concentration of the chemicals. (6 marks)

Stage 1 : Improved insulation

• Insulate the beaker or use a polystyrene cup or a lid

• To reduce heat loss

Stage 2 : Improved temperature recording

• Record the temperature for a suitable time before adding the metal

• To establish an accurate initial temperature

• Record temperature values at regular time intervals

• To plot the temperature results against time on a graph

Stage 3 : Improved analysis of results

• Extrapolate the cooling back to the point of addition

• To establish a (theoretical) maximum temperature OR temperature change

Explain why the enthalpy of hydration becomes less exothermic from Li+ to K+. (2 marks)

• (From Li+ to K+ ) size (of ion) increases

• Electrostatic attraction between metal ion and O^δ– of water decreases

Explain why the enthalpy of lattice dissociation for sodium oxide is greater than the enthalpy of lattice dissociation for sodium chloride. (2 marks)

• Oxide ions have higher negative charge OR smaller size OR higher charge density/higher charge/size ratio than chloride ions

• stronger attraction between O^2- and Na⁺ ions

Give a reason why data books do not contain a value for the enthalpy of solution of sodium oxide. (1 mark)

It reacts with water OR It reacts to form a solution of NaOH

Explain why the enthalpy of hydration of fluoride ions is more negative than the enthalpy of hydration of chloride ions. (2 marks)

• Fluoride ions/F⁻ ions are smaller OR F⁻ has a higher charge density

• Stronger attraction of fluoride ion to δ+ on H/ electron deficient H in water

Table 3 shows lattice enthalpies based on a perfect ionic model and lattice enthalpies from Born–Haber cycles for three metal chlorides.

Discuss the values in Table 3. In your answer you should

• compare the three values based on a perfect ionic model

• compare the values based on a perfect ionic model to the values from a Born–Haber cycle for each compound.

(6 marks)

Stage 1: comparing values from perfect ionic model

• Value for CaCl₂ is larger OR Values for KCl and AgCl are similar OR Values for CaCl₂ > AgCl > KCl

• Ca²⁺ has a larger charge/ is a smaller ion OR Ag⁺ and K⁺ have smaller charge or larger ions

• CaCl₂ has stronger ionic bonds or stronger attraction between + and - ions (Ca²⁺and Cl^- ) OR AgCl and KCl have weaker ionic bonds or weaker attraction between + and - ions (Ag⁺/ K⁺ and Cl^- )

Stage 2: similarities in the perfect ionic model and Born-Haber cycle

• CaCl₂ has similar values (between the perfect ionic model and Born-Haber cycle)

• KCl has similar values (between the perfect ionic model and Born-Haber cycle)

• CaCl2 and KCl have (almost) perfect ionic bonding or + ions are point charges/(perfectly) spherical

Stage 3 difference in the perfect ionic model and Born-Haber cycle values

• AgCl has larger difference in values (between the perfect ionic model and Born-Haber cycle)

• AgCl contains (some) covalent character

• Ag⁺ more polarising/distorts electron cloud more

Suggest a change to the student’s method, using the same apparatus, that would reduce the percentage uncertainty in the temperature change. Give a reason for your answer. (2 marks)

• use a larger mass/amount of NH₄NO₃ / solid

• so temperature change/decrease is greater OR final temperature is lower