Solid ammonium dichromate decomposes on heating according to the equation:

$(\text{NH}_4)_2\text{Cr}_2O_7(s) \to \text{Cr}_2O_3(s) + N_2(g) + 4H_2O(g)$

The third law of thermodynamics is often described using a perfect crystal at $0\ \text{K}$.

The Haber process is represented by the equation:

$N_2(g) + 3H_2(g) \to 2NH_3(g)$

The standard molar entropies are: $S^\circ[N_2(g)] = 192\ \text{J K}^{-1}\text{ mol}^{-1}$, $S^\circ[H_2(g)] = 131\ \text{J K}^{-1}\text{ mol}^{-1}$ and $S^\circ[NH_3(g)] = 193\ \text{J K}^{-1}\text{ mol}^{-1}$.

For a reaction at $298\ \text{K}$, $\Delta H^\circ = +82.4\ \text{kJ mol}^{-1}$ and $\Delta S^\circ = +172\ \text{J K}^{-1}\text{ mol}^{-1}$.

Phosgene formation is represented by:

$CO(g) + Cl_2(g) \rightleftharpoons COCl_2(g)$

At a particular temperature, $K = 4.0 \times 10^3$. For a reacting mixture at the same temperature, $Q = 2.0 \times 10^5$.

A voltaic cell has $E^\circ_{\text{cell}} = +0.460\ \text{V}$ for its overall reaction. Two moles of electrons are transferred per mole of reaction.

Ammonium nitrate dissolves readily in water, although the process is endothermic.

The table shows standard molar entropy values, $S^\circ$, for selected substances at $298\ \text{K}$.

The graph shows the Gibbs energy change, $\Delta G$, for the forward reaction as a reversible reaction mixture changes composition at constant temperature.

The decomposition of calcium carbonate is represented by:

$\text{CaCO}_3(s) \to \text{CaO}(s) + CO_2(g)$

For this reaction, $\Delta H^\circ = +178.3\ \text{kJ mol}^{-1}$ and $\Delta S^\circ = +160.5\ \text{J K}^{-1}\text{ mol}^{-1}$.

Two reactions have the following thermodynamic signs.

Reaction I: $\Delta H^\circ < 0$ and $\Delta S^\circ < 0$

Reaction II: $\Delta H^\circ > 0$ and $\Delta S^\circ > 0$

A graph of $\Delta G^\circ$ against temperature is shown for a reaction.

At $298\ \text{K}$, a reaction has $\Delta G^\circ = -5.70\ \text{kJ mol}^{-1}$.

The table gives standard molar entropy values for the substances in the Haber process.

$\text{N}_2(g) + 3\text{H}_2(g) \to 2\text{NH}_3(g)$

The thermal decomposition of calcium carbonate is represented by the equation.

$\text{CaCO}_3(s) \to \text{CaO}(s) + \text{CO}_2(g)$

Thermodynamic data for the substances are shown.

The table gives $\Delta H^\circ$ and $\Delta S^\circ$ values for four reactions.

A pure liquid $X$ is heated at standard pressure. The graph shows $\Delta G$ for the process $X(l) \to X(g)$ as temperature changes. The enthalpy change of vaporization is also shown.

Electrochemical data for three standard cells are shown. The relationship between standard Gibbs energy change and standard cell potential is $\Delta G^\circ = -nFE^\circ_{\text{cell}}$.

For a reaction at $298\ \text{K}$, $\Delta G^\circ = +12.0\ \text{kJ mol}^{-1}$. At one instant the reaction quotient is $Q = 0.010$.

The graph shows how $\Delta G^\circ$ varies with temperature for a reaction under standard conditions.

The table shows thermodynamic data for four reactions under standard conditions.

Ammonium nitrate dissolves endothermically in water. The table gives thermodynamic data for the dissolution process at $298\ \text{K}$.

$\text{NH}_4\text{NO}_3(s) \to \text{NH}_4^+(aq) + \text{NO}_3^-(aq)$

Ethanol can exist as a solid, liquid or gas. A perfect crystal of ethanol is considered at very low temperature. The enthalpy of vaporization of ethanol at its normal boiling point, $351\ \text{K}$, is $38.6\ \text{kJ mol}^{-1}$.

Methanol can be synthesized by the reaction.

$\text{CO}(g) + 2\text{H}_2(g) \rightleftharpoons \text{CH}_3\text{OH}(g)$

The table gives the composition of a reaction mixture at $500\ \text{K}$ and the value of $K$ at this temperature.

The reaction $\text{N}_2\text{O}_4(g) \rightleftharpoons 2\text{NO}_2(g)$ has a positive standard Gibbs energy change at $298\ \text{K}$. Data for a non-equilibrium mixture are shown.

Calcium carbonate decomposes on heating according to the equation:

$\text{CaCO}_3(s) \to \text{CaO}(s) + CO_2(g)$

For this reaction, $\Delta H^\circ = +178.3\ \text{kJ mol}^{-1}$. The standard entropy values are $S^\circ(\text{CaCO}_3(s)) = 92.9\ \text{J K}^{-1}\text{ mol}^{-1}$, $S^\circ(\text{CaO}(s)) = 39.8\ \text{J K}^{-1}\text{ mol}^{-1}$ and $S^\circ(CO_2(g)) = 213.8\ \text{J K}^{-1}\text{ mol}^{-1}$.

Ammonia is formed in the Haber process:

$N_2(g)+3H_2(g) \rightleftharpoons 2NH_3(g)$

For the forward reaction, $\Delta H^\circ = -92.4\ \text{kJ mol}^{-1}$. Use $S^\circ(N_2(g))=191.5\ \text{J K}^{-1}\text{ mol}^{-1}$, $S^\circ(H_2(g))=130.7\ \text{J K}^{-1}\text{ mol}^{-1}$ and $S^\circ(NH_3(g))=192.8\ \text{J K}^{-1}\text{ mol}^{-1}$.

Ammonium nitrate dissolves endothermically in water and is used in some cold packs:

$NH_4NO_3(s) \to NH_4^+(aq)+NO_3^-(aq)$

For this process, $\Delta H^\circ = +25.7\ \text{kJ mol}^{-1}$ and $\Delta S^\circ_{\text{system}} = +108.7\ \text{J K}^{-1}\text{ mol}^{-1}$.

The graph shows how $\Delta G^\circ$ varies with temperature for two reactions, A and B. The straight-line relationships are:

A: $\Delta G^\circ = 84.0 - 0.120T$

B: $\Delta G^\circ = -50.0 + 0.080T$

$\Delta G^\circ$ is in $\text{kJ mol}^{-1}$ and $T$ is in $\text{K}$.