The chromate-dichromate equilibrium is shown.

$2CrO_4^{2-}(aq) + 2H^+(aq) \rightleftharpoons Cr_2O_7^{2-}(aq) + H_2O(l)$

yellow $CrO_4^{2-}$, orange $Cr_2O_7^{2-}$

A few drops of aqueous sodium hydroxide are added to an equilibrium mixture at constant temperature.

In a sealed bottle of carbonated water, carbon dioxide is involved in the heterogeneous equilibrium:

$CO_2(g) \rightleftharpoons CO_2(aq)$

The bottle is opened at constant temperature.

Dinitrogen tetroxide and nitrogen dioxide form the following equilibrium in a sealed syringe.

$N_2O_4(g) \rightleftharpoons 2NO_2(g)$

The forward reaction is endothermic. $N_2O_4$ is colourless and $NO_2$ is brown.

At a particular temperature, $K = 49.0$ for the equilibrium:

$H_2(g) + I_2(g) \rightleftharpoons 2HI(g)$

At one instant, the concentrations are:

$[H_2] = 0.0200\ \text{mol dm}^{-3}$, $[I_2] = 0.0100\ \text{mol dm}^{-3}$ and $[HI] = 0.0600\ \text{mol dm}^{-3}$.

A homogeneous equilibrium is established at constant temperature.

$A(g) + B(g) \rightleftharpoons C(g)$

Initially, $[A] = 0.500\ \text{mol dm}^{-3}$ and $[B] = 0.500\ \text{mol dm}^{-3}$, with no $C$ present. At equilibrium, $[C] = 0.300\ \text{mol dm}^{-3}$.

A small amount of liquid bromine is placed in a sealed glass vessel at constant temperature. The graph shows the rates of evaporation and condensation of bromine over time.

A $0.200\ \text{mol dm}^{-3}$ solution of a weak acid, $HA$, is prepared at $298\ \text{K}$.

$HA(aq) \rightleftharpoons H^+(aq) + A^-(aq)$

$K_a = 1.6 \times 10^{-5}$

Use the approximation $[HA]_{\text{initial}} \approx [HA]_{\text{eqm}}$.

For a reaction at $298\ \text{K}$, the standard Gibbs energy change is $+12.4\ \text{kJ mol}^{-1}$.

Use $\Delta G^\circ = -RT\ln K$ and $R = 8.31\ \text{J K}^{-1}\ \text{mol}^{-1}$.

At $350\ \text{K}$, a homogeneous equilibrium has $K = 2.50 \times 10^4$.

Use $\Delta G^\circ = -RT\ln K$ and $R = 8.31\ \text{J K}^{-1}\ \text{mol}^{-1}$.

Nitrosyl chloride decomposes in a sealed vessel according to the homogeneous equilibrium:

$2\text{NOCl}(g) \rightleftharpoons 2\text{NO}(g) + \text{Cl}_2(g)$

The table gives equilibrium concentrations at one temperature.

The equilibrium constant for the formation of methanol is shown at different temperatures.

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

The equilibrium below is established in aqueous solution. The product ion is red.

$\text{Fe}^{3+}(aq)+\text{SCN}^-(aq) \rightleftharpoons \text{FeSCN}^{2+}(aq)$

Absorbance is used to monitor the concentration of $\text{FeSCN}^{2+}$ after different additions.

At a fixed temperature, $K=50.0$ for the equilibrium:

$\text{H}_2(g)+\text{I}_2(g) \rightleftharpoons 2\text{HI}(g)$

The table gives concentrations in two reaction mixtures.

The equilibrium constant for the reaction below is $K = 0.250$ at a particular temperature.

$2A(g) \rightleftharpoons B(g)$

At one instant, $[A] = 0.400\ \text{mol dm}^{-3}$ and $[B] = 0.0200\ \text{mol dm}^{-3}$.

Carbon dioxide is allowed to reach equilibrium between the gas phase and water in a sealed syringe.

$\text{CO}_2(g) \rightleftharpoons \text{CO}_2(aq)$

The plunger is then pushed in at constant temperature.

Dinitrogen tetroxide and nitrogen dioxide form the equilibrium below. Nitrogen dioxide is brown and dinitrogen tetroxide is colourless.

$\text{N}_2\text{O}_4(g) \rightleftharpoons 2\text{NO}_2(g)$

For the reaction as written, $\Delta H_r^\circ=+57\ \text{kJ mol}^{-1}$. The diagrams show the colour intensity of sealed tubes under different conditions.

A homogeneous equilibrium is established at constant temperature.

$A(g)+B(g) \rightleftharpoons C(g)$

The table shows initial concentrations and one measured equilibrium concentration.

A weak monoprotic acid, HA, is dissolved in water.

$\text{HA}(aq) \rightleftharpoons \text{H}^+(aq)+\text{A}^-(aq)$

At $298\ \text{K}$, $K_a=1.8\times10^{-5}$ for HA. The initial concentration of HA is $0.100\ \text{mol dm}^{-3}$.

The equilibrium constant for a reaction at $298\ \text{K}$ is measured as $K=3.2\times10^4$.

The relationship between standard Gibbs energy change and $K$ is:

$\Delta G^\circ=-RT\ln K$

A small amount of liquid bromine is placed in a sealed flask. After some time the intensity of the brown vapour becomes constant.

Methanol can be manufactured from carbon monoxide and hydrogen in a reversible reaction.

$CO(g)+2H_2(g)\rightleftharpoons CH_3OH(g)$

Ammonia is produced in the Haber process.

$N_2(g)+3H_2(g)\rightleftharpoons 2NH_3(g)\qquad \Delta H<0$

For a reversible reaction at constant temperature, the reaction quotient $Q$ was calculated at different times after mixing the reactants. The equilibrium constant is $K=12$ at this temperature.

For the equilibrium below at $298\ \text{K}$, the standard Gibbs energy change is $+4.80\ \text{kJ mol}^{-1}$.

$\text{N}_2\text{O}_4(g) \rightleftharpoons 2\text{NO}_2(g)$

A mixture is then prepared with non-equilibrium concentrations of the two gases.

An aqueous equilibrium involving cobalt ions is shown.

$[\text{Co(H}_2\text{O)}_6]^{2+}(aq)+4Cl^-(aq)\rightleftharpoons [\text{CoCl}_4]^{2-}(aq)+6H_2O(l)$

The forward reaction is endothermic. The octahedral complex is pink and the tetrahedral complex is blue.

Nitrogen dioxide and dinitrogen tetroxide form an equilibrium in a sealed gas syringe.

$N_2O_4(g)\rightleftharpoons 2NO_2(g)$

$N_2O_4$ is colourless and $NO_2$ is brown.

A gas $X$ is used in an industrial absorption process represented by the heterogeneous equilibrium:

$X(g)\rightleftharpoons X(aq)$

The dissolution of $X$ is exothermic.

At $700\ \text{K}$, the equilibrium constant for the reaction below is $K=4.0$.

$CO(g)+H_2O(g)\rightleftharpoons CO_2(g)+H_2(g)$

At one instant, $[CO]=0.50\ \text{mol dm}^{-3}$, $[H_2O]=0.40\ \text{mol dm}^{-3}$, $[CO_2]=0.60\ \text{mol dm}^{-3}$ and $[H_2]=0.20\ \text{mol dm}^{-3}$.

Methanoic acid is a weak acid. At $298\ \text{K}$, $K_a=1.8\times10^{-4}$ for:

$HCOOH(aq)\rightleftharpoons H^+(aq)+HCOO^-(aq)$

A solution initially contains $0.100\ \text{mol dm}^{-3}$ methanoic acid.

At $298\ \text{K}$, the equilibrium constant for a reaction is $K=3.2\times10^{-5}$. The standard Gibbs energy change is related to $K$ by

$\Delta G^\circ=-RT\ln K$

At a particular temperature, $K=1.80$ for the equilibrium:

$A(g)+B(g)\rightleftharpoons 2C(g)$

A mixture initially contains $0.600\ \text{mol dm}^{-3}$ of $A$, $0.600\ \text{mol dm}^{-3}$ of $B$ and no $C$.