A-Level Chemistry: Entropy, Lattice Energy, and Feasibility

Flashcards covering entropy, spontaneity, Born-Haber cycles, lattice energies, and enthalpy of solution.

What is a spontaneous process?

A process that proceeds on its own without any external influence (for example, diffusion).

How is entropy ($S$) defined in relation to energy distribution?

Entropy is a description of the number of ways atoms can share quanta of energy ($W$). When the number of ways of arranging the energy is high, the system is disordered and entropy is high.

Why is there a larger jump in entropy during boiling compared to melting?

Gases are much more disordered than liquids and have significantly more ways of arranging energy quanta, resulting in a much larger entropy increase during boiling.

What is the entropy of a substance at absolute zero ($0\,\text{K}$)?

At $0\,\text{K}$, substances have zero entropy because only perfect crystals exist with stationary particles and no disorder.

What two qualitative changes generally cause a significant increase in the entropy of a system ($\Delta S^\circ_{\text{system}}$)?

1. A change of state from solid or liquid to gas.\n2. A significant increase in the number of molecules between products and reactants.

What is the equation used to calculate the standard entropy change of a system ($\Delta S^\circ_{\text{system}}$)?

$\Delta S^\circ_{\text{system}} = \sum S^\circ_{\text{products}} - \sum S^\circ_{\text{reactants}}$ with units of $\text{J}\,\text{K}^{-1}\,\text{mol}^{-1}$.

How is the entropy change of the surroundings ($\Delta S_{\text{surrounding}}$) calculated?

$\Delta S_{\text{surrounding}} = -\frac{\Delta H_{\text{reaction}}}{T}$, where $\Delta H_{\text{reaction}}$ is converted to $\text{J}\,\text{mol}^{-1}$ (multiplied by $1000$) and $T$ is temperature in $\text{K}$.

What condition must be met for a reaction to be spontaneous regarding $\Delta S_{\text{total}}$?

$\Delta S_{\text{total}}$ must be positive ($\Delta S_{\text{total}} > 0$). If $\Delta S_{\text{total}}$ is negative, the reaction is not feasible.

How is the temperature calculated at which a reaction becomes feasible?

By setting $\Delta S_{\text{total}} = 0$ in the equation $0 = -\frac{\Delta H}{T} + \Delta S_{\text{system}}$, giving $T = \frac{\Delta H}{\Delta S_{\text{system}}}$.

Why might a reaction with a negative free energy change ($\Delta G < 0$) still fail to occur?

The reaction may have a high activation energy ($E_A$), meaning kinetic factors limit the reaction rate so it occurs too slowly to be observed.

What is the standard enthalpy change of formation ($\Delta_f H$)?

The energy transferred when $1\,\text{mole}$ of a compound is formed from its elements under standard conditions ($298\,\text{K}$ and $100\,\text{kPa}$), with all reactants and products in their standard states.

What is the enthalpy of atomisation ($\Delta_{\text{at}} H$)?

The enthalpy change when $1\,\text{mole}$ of gaseous atoms is formed from an element in its standard state.

Why is the second electron affinity of oxygen endothermic?

It takes energy to overcome the electrostatic repulsive force between the incoming negative electron and the already negatively charged $\text{O}^-$ gaseous ion.

What is the standard enthalpy of lattice formation ($\Delta_{\text{LE}} H$)?

The standard enthalpy change when $1\,\text{mole}$ of an ionic crystal lattice is formed from its constituent ions in gaseous form.

How do ionic radius and ionic charge affect lattice formation enthalpy?

Smaller ions and higher ionic charges increase the electrostatic attraction between ions, making the lattice formation enthalpy more negative (stronger lattice).

What assumptions are made by the perfect ionic model?

The ions are assumed to be $100\%$ ionic, perfectly spherical, and held together solely by purely electrostatic attractions.

Why do experimental Born-Haber lattice enthalpies have a larger magnitude than theoretical values for certain compounds?

Covalent character causes polarization (distortion of the anion charge cloud), creating partial covalent bonding that strengthens the lattice beyond purely ionic electrostatic attraction.

What two factors increase the polarising power of a cation?

A small ionic radius and a high positive charge (multiple charges).

Why is $\text{CaCl}_2$ formed preferentially over $\text{CaCl}$ and $\text{CaCl}_3$?

$\text{CaCl}_2$ has the most exothermic enthalpy of formation ($\Delta_f H = -739.2\,\text{kJ}\,\text{mol}^{-1}$) because the stronger lattice enthalpy more than compensates for the second ionization energy, whereas forming $\text{CaCl}_3$ requires too much energy to remove the third electron.

What is the enthalpy of hydration ($\Delta_{\text{hyd}} H$)?

The enthalpy change when $1\,\text{mole}$ of gaseous ions becomes hydrated such that further dilution causes no further heat change; it is always exothermic.

What is the formula relating enthalpy of solution ($\Delta_{\text{sol}} H$), lattice formation enthalpy ($\Delta_{\text{LE}} H$), and hydration enthalpy ($\Delta_{\text{hyd}} H$)?

$$\Delta_{\text{sol}} H = -\Delta_{\text{LE}} H + \sum \Delta_{\text{hyd}} H$$

Why can an ionic salt dissolve even if its enthalpy of solution ($\Delta_{\text{sol}} H$) is endothermic?

Dissolving increases disorder, giving a positive $\Delta S_{\text{system}}$, which can outweigh the negative $-\frac{\Delta H}{T}$ term to make $\Delta S_{\text{total}}$ positive, especially at higher temperatures.