Reversible reaction: A chemical reaction that can occur in both the forward and reverse direction and shown by two arrows .
It is represented in the general case by:
aA + bB cC + dD
The expression for Q may be written down by inspection of the balanced chemical equation. It is the product of the instantaneous concentrations of the reaction products each raised to the power of its stoichiometric coefficent divided by the product of the instantaneous concentrations of the reactants each raised to the power of its stoichiometric coefficient. Here the square bracket notation, e.g. [A], is used as the symbol for concentration rather than c(A). The form c(A), usually represents the overall concentration of a species disregarding any reactions which might have occurred. (See below, equilibrium constant.)
Chemical equilibrium: The state of a reversible chemical reaction when the rate of the reaction in the reverse direction equals that in the forward direction; there is no change in amounts or concentrations of reactants or products with time. As the reaction is occurring in both directions we called it dynamic equilibrium. It requires an input of energy to change the composition at equilibrium, i.e. to drive the reaction in the forward or reverse direction. A reaction at equilibrium is indicated by the use of two half arrows, , as shown for the general case:
aA + bB cC + dD
Equilibrium constant: Symbol K, a quantitative expression for the equilibrium composition of a chemical reaction. It is the value of the reaction quotient, Q, at equilibrium. Thus for the general reaction above
where [A] etc. are the equilibrium concentrations of the reactant and product species present. Like Q the expression for K may be written down by inspection of the balanced chemical equation. Again the square bracket notation is used for concentration.
Pressure equilibrium constant: Symbol Kp, an equilibrium constant in which the equilibrium composition is expressed in terms of partial pressures of reactants and products instead of concentrations; because the concentration of a gas is directly proportional to its pressure (partial pressure in a mixture of gases).
Standard (reference) concentrations and pressures: Standard or reference concentrations and pressures are useful concepts.
Units of equilibrium constants, K: In principle, the magnitude and units of an equilibrium constant will depend on the stoichiometry of the reaction and on the chosen units of concentration. To avoid having different values for the same equilibrium constant chemists use the ratio of the concentration to a standard concentration, or ratio of the partial pressure to a standard pressure, and K is dimensionless, i.e. just a number without units.
Thus where the values of the concentrations of the solute species are just the numerical values of their concentrations in mol L¯1.
Position of equilibrium: A qualitative expression for the relative amounts of reactants and products present at equilibrium. The position of equilibrium may lie well to the left with only very small but measurable amounts of products present, or well to the right with only small but measurable amounts of reactants present, or anywhere in between. The magnitude of K is a quantitative measure of the position of the equilibrium. If K is large, then the numerator is much larger than the denominator, the concentration of the products is much greater than that of the reactants, and the equilibrium lies well to the right. If K is small the equilibrium lies well to the left, and only a small fraction of reactants are converted into products.
pK: pK = −log10K. Many equilibrium constants have extremely small numerical values and it is often convenient, especially in tabulation, to express them in a logarithmic form.
Temperature control of equilibrium: The magnitude of K varies with temperature, decreasing with increasing temperature for exothermic reactions, and increasing with increasing temperature for endothermic reactions.
K for overall reactions: When an overall reaction can be considered as the sum of a series of reactions, K for the overall reaction is the product of the equilibrium constants of the individual reactions, as can be seen for a general case:
It follows from the above discussion (and also from the definition of K) that if the stoichiometry of a reaction equation is altered by multiplying by 2, the value of K is squared.