Thermodynamics (OCR A Level Chemistry A): Flashcards

False.

Electron affinity is exothermic, so the arrow points downwards in the diagram, representing a decrease in energy.

It is the direct route from elements in their standard states to the ionic lattice; for sodium chloride it is exothermic, so the arrow points downwards.

In a Born-Haber cycle, the enthalpy of formation provides the direct route, while all other steps (atomisation, ionisation, electron affinity, lattice enthalpy) form the indirect route.

Removing an electron from a gaseous sodium atom requires energy, making it an endothermic process, so it is represented by an upward arrow indicating an increase in energy level.

Hess's Law states that the total enthalpy change for a reaction is independent of the route taken, provided the initial and final conditions are the same.

By applying Hess's Law: ΔH^ꝋ_latt = ΔH^ꝋ_f - ΔH^ꝋ₁, where ΔH^ꝋ₁ is the sum of all enthalpy changes needed to convert elements in their standard states to gaseous ions.

In a Born-Haber cycle calculation, the stage you are asked to find is always the direct route, and all other known stages form the indirect route.

True.

MgCl₂ contains two moles of Cl^- ions, so 2 moles of electrons are added to 2 moles of chlorine atoms, requiring the electron affinity value to be multiplied by 2.

Because the stoichiometry of the ionic compound determines how many moles of each ion are present; for example, MgCl₂ requires two moles of Cl^-, so the electron affinity must be multiplied by 2.

The simplified Born-Haber equation is: ΔH^ꝋ_f = ΔH^ꝋ₁ + ΔH^ꝋ_latt, where ΔH^ꝋ₁ represents the sum of all steps converting elements in their standard states to gaseous ions.

It is the sum of all intermediate enthalpy changes needed to convert elements in their standard states into gaseous ions, including atomisation, ionisation energies and electron affinities.

An enthalpy of solution 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.

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

An enthalpy of hydration is the enthalpy change when 1 mole of a specified gaseous ion dissolves in sufficient water to form an infinitely dilute solution. It is always exothermic.

When an ionic solid dissolves, the oxygen atom of water is attracted to positive ions, while the hydrogen atoms are attracted to negative ions.

False.

The enthalpy of solution can be either exothermic or endothermic, depending on the relative magnitudes of the lattice enthalpy and the hydration enthalpies of the ions.

By Hess's Law: ΔH^ꝋ_hyd = ΔH^ꝋ_latt + ΔH^ꝋ_sol, where the total hydration enthalpy is the sum of the hydration enthalpies of all cations and anions.

To calculate ΔH^ꝋ_hyd of one ion, you need the lattice enthalpy, the enthalpy of solution and the hydration enthalpy of the other ion.

MgCl₂ contains two moles of Cl^- ions per formula unit, so the total hydration enthalpy from chloride ions must account for 2 moles of Cl^-.

The charge density of an ion is the ratio of its ionic charge to its volume. Ions with a higher charge density exert stronger electrostatic forces on surrounding ions or polar molecules.

Larger ions have a lower charge density, so the electrostatic forces of attraction between oppositely charged ions in the lattice are weaker. Less energy is released when the lattice forms, making the lattice enthalpy less exothermic.

The lattice enthalpy of CsF is less exothermic than that of KF because the Cs⁺ ion is larger than the K⁺ ion.

True.

Greater ionic charge produces higher charge density, leading to stronger electrostatic attractions between ions in the lattice and more energy released on formation.

Ca²⁺ and O^2- ions carry a greater ionic charge than K⁺ and Cl^-, giving them a higher charge density and stronger electrostatic attractions in the lattice. Ca²⁺ and O^2- are also smaller ions, which further increases the attraction.

Smaller ions have a greater charge density, producing stronger ion-dipole attractions between the ion and surrounding water molecules. More energy is released as the ions become hydrated.

Ions with a larger ionic charge have a greater charge density, resulting in a more exothermic standard enthalpy of hydration.

Entropy (S) is a measure of the number of possible arrangements of particles and their energy in a system, reflecting the degree of disorder or chaos. A more disordered system has a higher entropy and is energetically more stable.

Enthalpy change only accounts for energy stability, but does not explain why a system would move to a less stable, higher-energy state. The entropy of the system must also be considered, as an increase in disorder can drive a reaction.

True.

Particles in a liquid are arranged more randomly and can move past each other, making the system more disordered than the fixed positions found in a solid.

Entropy is measured in joules rather than kilojoules because entropy changes are an order of magnitude smaller than enthalpy changes.

The reaction produces CO₂ gas, which is far more disordered than the solid reactant CaCO₃ because gas molecules move freely and randomly. The system therefore moves to a state of greater disorder, increasing entropy.

The standard entropy change is calculated using ΔS^ꝋ = ΣS^ꝋ_products − ΣS^ꝋ_reactants, where Σ means the sum of standard entropies for each species. Unlike enthalpy of formation, the standard entropy of elements is not zero.

The standard entropy change is found by subtracting the sum of reactant entropies from the sum of product entropies.

False.

Feasibility only indicates that a reaction is energetically favourable; it takes no account of the rate of reaction, so a feasible reaction may be extremely slow.

A feasible reaction (or spontaneous reaction) is one that is energetically favourable and will occur of its own accord. Feasibility is determined by the sign of the Gibbs free energy change, ΔG.

A negative ΔG means the reaction is feasible and likely to occur spontaneously. The system releases free energy, making the products thermodynamically more stable than the reactants.

When ΔG is negative, the reaction is feasible; when ΔG is positive, the reaction is not feasible.

True.

When ΔH is positive and ΔS is negative, both terms in the Gibbs equation contribute positively to ΔG, so ΔG is always positive regardless of temperature.

At high temperatures, the −TΔS term becomes large and negative enough to overcome the positive ΔH, making ΔG negative and the reaction feasible. At low temperatures, the −TΔS term is too small to offset the positive ΔH.

When ΔH is negative and ΔS is positive, both the ΔH term and the −TΔS term are negative, so ΔG is always negative at any temperature. The reaction is therefore always thermodynamically feasible.

True.

ΔH is expressed in kJ mol^-1, so ΔS must be divided by 1000 to give consistent units before calculating ΔG.

Gibbs free energy (G) is a thermodynamic quantity that combines enthalpy and entropy to determine whether a reaction is feasible. A reaction is feasible when the change in Gibbs free energy, ΔG, is negative.

The first method uses the Gibbs equation, ΔG^ꝋ = ΔH^ꝋ − TΔS^ꝋ, with separately known enthalpy and entropy values. The second method uses ΔG^ꝋ = ΣΔG^ꝋ_products − ΣΔG^ꝋ_reactants, applying standard free energy of formation data.

At equilibrium, the Gibbs free energy change equals zero, representing the minimum energy point of the system.

False.

ΔG only tells you that a reaction is thermodynamically feasible; it gives no information about kinetics or how quickly the reaction proceeds.

Although the reaction is thermodynamically feasible (ΔG = −117 kJ mol^-1), it has a high activation energy that prevents it from proceeding at a measurable rate. A catalyst such as MnO₂ is needed to provide an alternative pathway with lower activation energy.

When ΔG is negative, ln K is positive, meaning the equilibrium constant K is greater than 1. This shows that the products are favoured at equilibrium.

When using ΔS values in the Gibbs equation, the units must be divided by 1000 to convert from J K^-1 mol^-1 to kJ K^-1 mol^-1.

True.

For elements in their standard state, the standard free energy of formation is defined as zero, just as with standard enthalpy of formation.