How far? (5.1.2)

Homogeneous equilibria involve reactants and products that are in the same phase.

The equilibrium constant, , is the reaction quotient at equilibria, and represents the ratio of the concentrations of products to reactants at equilibrium, in homogeneous reactions.

, , , denote the concentrations of the respective substances at equilibrium, and , , , are the stoichiometric coefficients of the balanced equation.

Calculate by inputting the equilibrium concentrations of each species into the expression for .

Determine units for by inserting the unit for concentration into the expression and applying the laws of indices.

Heterogeneous equilibria involve reactants and products in different phases.

In these equilibria, the concentration of pure solids and liquids are considered constant and therefore not included in the equilibrium expressions.

For this heterogeneous system, the equilibrium constant, K_c, is given by:

Where:

• and are the concentrations of each species in
• and are the stoichiometric coefficients from the balanced chemical equation.

and are solids and do not appear in the expression.

is the equilibrium constant for gases, expressed in terms of partial pressures.

It quantifies the ratio of gaseous products and reactants at equilibrium.

In a gaseous chemical system, each gas exerts its own partial pressure. The sum of these partial pressures is equal to the total pressure of the system.

To determine the partial pressure of a gas, you multiply its mole fraction by the total system pressure:

The mole fraction of a component in a mixture is defined as the ratio of the number of moles of that component to the total number of moles of all components in the mixture.

It can be calculated using:

For a gaseous equilibrium:

The equilibrium constant in terms of partial pressures, is given by:

Where:

• is the equilibrium partial pressure of species
• , , , and are the stoichiometric coefficients from the balanced chemical equation.

Partial pressures can be in any pressure unit but must align across the equation.

The units of are established by inputting the units for partial pressure into the expression.

Heterogeneous equilibria involve reactants and products in different phases. In these equilibria, solids and pure liquids (solvents) are not included in the equilibrium expressions.

With the following homogeneous system:

the equilibrium constant, , only includes the gaseous reactants and is given by:

Where:

• is the equilibrium constant
• is the equilibrium partial pressure of each gaseous species
• and are the stoichiometric coefficients from the balanced chemical equation

and are solids and do not appear in the expression.

The most practical methods for determining equilibrium concentrations do not alter the chemical makeup of the reaction mixture. Methods include colorimetry, using a pH probe and other calibrated instruments.

To use titration to calculate equilibrium concentrations the reaction mixture must first be quenched.

Quenching (stopping) a chemical reaction prevents the introduction of another species via titration from altering the position of equilibria.

The equilibrium constants, and remain constant at a given temperature; when temperature changes, the equilibrium constant also changes.

Partial pressure or concentration changes cause equilibrium shifts to regain the ratio described by or respectively, but the equilibrium constant remains unchanged.

Catalysts increase the rate at which equilibrium is established but do not change the position of equilibria.

Catalysts increase the rate of reaction by providing an alternative reaction pathway with a lower activation energy.

As they increase the rate of both the forward and reverse reactions equally, the presence of a catalyst does not affect the position of equilibrium or the value of the equilibrium constant. Catalysts simply allow the system to reach equilibrium faster.

In general terms:

The Haber process is commonly used to demonstrate the behaviour of gaseous equilibria.

Equilibrium constant, , is given by the equation:

If the volume of the container decreases, all the partial pressures will initially increase.

In this instance the equation’s output will become lower than that described by ; the denominator has increased more than the numerator.

The system will shift towards the right hand side, to increase the value of the numerator and decrease that of the denominator, restoring the balance as dictated by ​.

In general terms:

Changing temperature changes the value of the equilibrium constants, and .

The impact of temperature on the equilibrium constant, and therefore the position of equilibria, varies depending on whether the forward reaction is exothermic (releases heat, ) or endothermic (absorbs heat ).

Where the forward reaction is exothermic, increasing the temperature will decrease the equilibrium constant.

This happens because, according to Le Chatelier’s principle, the system will shift to the left to counteract the added heat energy by favouring the endothermic reverse reaction, which absorbs heat energy from the surroundings.

This shift results in an increase in the concentration of reactants and a decrease in the concentration of products, aligning to the new lower value of the equilibrium constant,

The acid dissociation constant is a specific type of equilibrium constant that applies to the dissociation of weak acids in water. It measures the strength of an acid in terms of its ability to dissociate into protons () and a conjugate base ().

For a generic weak acid, , the dissociation in water is represented as:

Which is simplified to:

The equilibrium constant expression for this reaction () is:

All the rules for position of equilibria which apply for also apply for .

The ionic product of water, , refers to the equilibrium constant for the self-ionisation of water. It provides insight into the concentration of hydrogen ions and hydroxide ions in pure water.

The self-ionisation of water is represented as:

Or more commonly simplified to:

The equilibrium constant expression for this reaction () is:

At 25 °C (298 K), the value of is:

All the rules for position of equilibria which apply for also apply for . As a result, the value of varies with temperature as does the of pure water, which is not always .