Chemistry › Definitions › Chemical equilibria

Chemical equilibria definitions

Every term the lessons in this unit define, lesson by lesson, worded as the lessons use them. Each lesson title links to its notes.

25 terms

Dynamic equilibrium, Le Chatelier's principle and industrial compromise

Reversible reaction: A reaction that can proceed in both the forward and the reverse direction under the same conditions.
Dynamic equilibrium: The state of a reversible reaction in a closed system in which the forward and reverse reactions proceed at equal rates, so the concentrations of reactants and products remain constant.
Closed system: A system that no substance can enter or leave, although energy may be transferred to or from it.
Open system: A system that substances can enter or leave, so a reversible reaction in it cannot reach equilibrium.
Position of equilibrium: The relative proportions of reactants and products in a mixture at equilibrium.
Le Chatelier's principle: If a system at equilibrium is subjected to a change in conditions, the position of equilibrium shifts in the direction that opposes the change.
Catalyst: A substance that increases the rate of a reaction without being changed in chemical composition or amount, by providing an alternative reaction route with a lower activation energy.
Equilibrium yield: The proportion of the reactants converted into products once the reaction mixture has reached equilibrium.
Compromise conditions: Conditions chosen to give an acceptable equilibrium yield at an acceptable rate and cost, where the conditions for the highest yield would make the reaction too slow or the plant too expensive.

Equilibrium constant, Kc: The concentrations of the products at equilibrium, each raised to the power of its coefficient and multiplied together, divided by the same expression for the reactants, at a stated temperature.
Equilibrium concentration: The concentration of a substance, in mol dm⁻³, in a mixture that has reached equilibrium; it is what each term in Kc stands for.
Homogeneous equilibrium: An equilibrium in which all the reactants and products are in the same phase.
Heterogeneous equilibrium: An equilibrium in which the reactants and products are in more than one phase, such as a solid with a gas.
Phase: A part of a system with the same composition and properties throughout: a solid, a liquid, a gas or a solution.
ICE table: A table of the initial amount, the change in amount and the equilibrium amount of each species, with the changes in the ratio of the coefficients in the equation.
Quenching: Slowing a reaction almost to a stop, by cooling, dilution or removing a reactant or catalyst, so that the composition of a sample does not change while it is analysed.
Reaction quotient, Q: The ratio with the same form as Kc but using the concentrations at any moment; it equals Kc only when the mixture is at equilibrium.
Partition coefficient, Kpc: The ratio of the concentrations of a solute in two immiscible solvents at equilibrium, at a stated temperature.

Mole fraction: The amount in moles of one gas in a mixture divided by the total amount in moles of all the gases.
Partial pressure: The pressure that one gas in a mixture would exert if it alone occupied the whole volume; its mole fraction multiplied by the total pressure.
Total pressure: The pressure of a gas mixture, equal to the sum of the partial pressures of all the gases in it.
Kp: The equilibrium constant for a reaction between gases, written with the partial pressure of each gas at equilibrium in place of its concentration.
Degree of dissociation: The fraction of the molecules of a substance that have dissociated when equilibrium is reached, written α.
Reaction quotient, Qp: The Kp expression evaluated with the partial pressures at any moment; it equals Kp only when the mixture is at equilibrium.
Atmosphere, atm: A unit of pressure equal to 101.325 kPa, the mean pressure of the air at sea level.