Energetics (Edexcel International A Level (IAL) Chemistry): Flashcards

1. Enthalpy of atomisation of Na

2. Enthalpy of atomisation of Cl

3. First ionisation energy of Na

4. First electron affinity of Cl

The cycle is completed by the enthalpy of formation of NaCl and the lattice enthalpy.

In a Born-Haber cycle, energy increases going upwards.

This means endothermic steps (atomisation, ionisation energy) point upwards and exothermic steps (electron affinity, lattice enthalpy) point downwards.

Δ_lattH^ꝋ = Δ_fH^ꝋ − Δ₁H^ꝋ

Where Δ₁H^ꝋ is the sum of all enthalpy changes needed to convert elements in their standard states to gaseous ions (atomisation + ionisation energies + electron affinities).

MgCl₂ contains two Cl^- ions.

Two moles of electrons are added to two moles of chlorine atoms, so the first electron affinity value must be multiplied by 2 in the calculation.

Δ_lattH^ꝋ = Δ_fH^ꝋ − [(Δ_atH^ꝋ K) + (Δ_atH^ꝋ Cl) + (IE₁ K) + (EA₁ Cl)]

The first electron affinity of chlorine is the final step in converting Cl atoms to Cl^- ions.

If the lattice enthalpy is given, the cycle can be used to calculate the enthalpy of formation of the ionic compound.

Any single unknown in the cycle can be found by rearranging the Hess's law expression.

False.

The enthalpy of formation can be exothermic or endothermic. In a Born-Haber cycle, an endothermic enthalpy of formation is shown by an arrow pointing upwards from the elements, displaced to the right.

Δ_lattH^ꝋ = (−437) − [(+90) + (+122) + (+418) + (−349)]

Δ_lattH^ꝋ = (−437) − (+281) = −718 kJ mol^-1

Polarisation occurs when a cation distorts the electron cloud of an anion, pulling electron density towards itself.

This introduces covalent character into what would otherwise be a purely ionic bond.

The theoretical value assumes a perfect ionic model with 100% spherical ions and purely electrostatic attraction.

In reality, polarisation introduces covalent character, causing the experimental lattice enthalpy to be more exothermic than predicted.

Polarising power is the ability of a cation to distort the electron cloud of an anion.

It increases with smaller ionic radius and higher ionic charge, as these give a greater charge density.

Li⁺ has greater polarising power.

Li⁺ has a smaller ionic radius than K⁺, giving it a higher charge density. A higher charge density means a greater ability to distort the electron cloud of an anion.

The ionic radius of the anion determines its polarisability.

Larger anions are more easily polarised as their outer electrons are further from the nucleus and less strongly attracted, making the electron cloud more easily distorted.

True.

Mg²⁺ has a greater charge and a smaller ionic radius than Na⁺, giving it a higher charge density and therefore greater polarising power.

Charge density is approximately proportional to charge / r²

Where r is the ionic radius. A higher charge or smaller radius gives a greater charge density and more polarising power.

Mg²⁺ has a higher charge and smaller radius than Na⁺, giving it a greater polarising power.

This means Mg²⁺ more strongly distorts the electron cloud of the Cl^- ion, introducing more covalent character.

Enthalpy of solution (Δ_solH^ꝋ) is the enthalpy change when 1 mole of an ionic substance dissolves in sufficient water to form an infinitely dilute solution.

It can be exothermic or endothermic.

Enthalpy of hydration (Δ_hydH^ꝋ) is the enthalpy change when 1 mole of a specified gaseous ion dissolves in sufficient water to form an infinitely dilute solution.

Hydration enthalpies are always exothermic.

Energy is released when ion-dipole attractions form between the gaseous ions and polar water molecules.

The δ− oxygen of water attracts positive ions and the δ+ hydrogen atoms attract negative ions.

Mg²⁺ (g) + aq → Mg²⁺ (aq)

This represents 1 mole of gaseous Mg²⁺ ions dissolving in excess water to form a hydrated ion.

False.

The enthalpy of solution (Δ_solH^ꝋ) can be either exothermic (negative) or endothermic (positive), depending on the relative magnitudes of the lattice enthalpy and hydration enthalpies.

1. The ionic lattice breaks down — this is endothermic (energy input required)

2. The ions become hydrated by water molecules — this is exothermic (energy released)

The overall Δ_solH^ꝋ depends on which process dominates.

KCl (s) + aq → K⁺ (aq) + Cl^- (aq)

One mole of solid KCl dissolves in sufficient water to form an infinitely dilute solution of its ions.

Δ_hydH^ꝋ = Δ_lattH^ꝋ + Δ_solH^ꝋ

This follows from Hess's law: the direct route (gaseous ions → aqueous ions) equals the indirect route (gaseous ions → ionic solid → aqueous ions).

To find the total enthalpy of hydration, the Δ_hydH^ꝋ values of the cation and anion must be added together.

Each ion hydrates independently, so both contributions must be summed.

Δ_hydH^ꝋ[total] = Δ_lattH^ꝋ + Δ_solH^ꝋ = (−711) + (+26) = −685 kJ mol^-1

Δ_hydH^ꝋ[Cl^-] = (−685) − (−322) = −363 kJ mol^-1

1. The entropy of the system (ΔS_system)

2. The enthalpy change of solution (Δ_solH^ꝋ)

3. The temperature of the water (in K)

Solubility is thermodynamically spontaneous when ΔS_total is positive.

False.

An endothermic Δ_solH^ꝋ can still allow dissolution if the entropy increase of the system is large enough to make ΔS_total positive, as seen with NH₄NO₃ dissolving in water.

MgCl₂ contains two Cl^- ions per formula unit.

Therefore 2 moles of Cl^- ions are hydrated per mole of MgCl₂, so Δ_hydH^ꝋ[Cl^-] must be multiplied by 2.

The entropy of water decreases.

Water molecules become more ordered as they arrange themselves around the cations and anions to form ion-dipole bonds.

ΔS_total = ΔS_system − Δ_solH / T

Where T is the temperature in kelvin. This shows that a positive ΔS_system and/or a negative Δ_solH (exothermic) both favour dissolution.

Increasing ionic radius makes lattice enthalpy less exothermic.\nLarger ions have more spread out charge and are further apart, weakening the electrostatic forces of attraction between oppositely charged ions.

Increasing ionic charge makes lattice enthalpy more exothermic.\nHigher charge gives greater charge density, resulting in stronger electrostatic attraction between ions in the lattice.

Ca²⁺ and O^2- have greater ionic charges than K⁺ and Cl^-.\nThe stronger electrostatic forces between the more highly charged ions release more energy when the lattice forms. Ca²⁺ and O^2- are also smaller ions, reinforcing the effect.

False.

Cs⁺ has a larger ionic radius than K⁺, giving weaker electrostatic attraction with F^-. The lattice enthalpy of CsF is therefore less exothermic than that of KF.

A higher ionic charge gives a greater charge density, which attracts water molecules more strongly.\nThis results in a more exothermic enthalpy of hydration, e.g. Mg²⁺ has a more negative Δ_hydH^ꝋ than Ba²⁺.

Smaller ions have higher charge density and attract water molecules more strongly.\nThis gives a more exothermic enthalpy of hydration, e.g. F^- has a more negative Δ_hydH^ꝋ than Cl^-.

A cation with a large charge and small radius has the greatest polarising power.

This combination gives the highest charge density, which most strongly distorts anion electron clouds and attracts water molecules.

True.

F^- has a smaller ionic radius than Cl^-, giving it a higher charge density. It attracts water molecules more strongly, releasing more energy on hydration.