EQUILIBRIUM

When the number of molecules leaving the liquid to vapour equals the number of

molecules returning to the liquid from vapour, equilibrium is said to be attained and is

dynamic in nature. Equilibrium can be established for both physical and chemical

processes and at this stage rate of forward and reverse reactions are equal. Equilibrium

constant, Kc is expressed as the concentration of products divided by reactants, each

term raised to the stoichiometric coefficient.

For reaction, a A + b B ƒ c C +d D

Kc = [C]c[D]d/[A]a[B]b

Equilibrium constant has constant value at a fixed temperature and at this stage

all the macroscopic properties such as concentration, pressure, etc. become constant.

For a gaseous reaction equilibrium constant is expressed as Kp and is written by replacing

concentration terms by partial pressures in Kc expression. The direction of reaction can

be predicted by reaction quotient Qc which is equal to Kc at equilibrium. Le Chatelier’s

principle states that the change in any factor such as temperature, pressure,

concentration, etc. will cause the equilibrium to shift in such a direction so as to reduce

or counteract the effect of the change. It can be used to study the effect of various

factors such as temperature, concentration, pressure, catalyst and inert gases on the

direction of equilibrium and to control the yield of products by controlling these factors.

Catalyst does not effect the equilibrium composition of a reaction mixture but increases

the rate of chemical reaction by making available a new lower energy pathway for

conversion of reactants to products and vice-versa.

All substances that conduct electricity in aqueous solutions are called electrolytes.

Acids, bases and salts are electrolytes and the conduction of electricity by their aqueous

solutions is due to anions and cations produced by the dissociation or ionization of

electrolytes in aqueous solution. The strong electrolytes are completely dissociated. In

weak electrolytes there is equilibrium between the ions and the unionized electrolyte

molecules. According to Arrhenius, acids give hydrogen ions while bases produce

hydroxyl ions in their aqueous solutions. Brönsted-Lowry on the other hand, defined

an acid as a proton donor and a base as a proton acceptor. When a Brönsted-Lowry

acid reacts with a base, it produces its conjugate base and a conjugate acid corresponding

to the base with which it reacts. Thus a conjugate pair of acid-base differs only by one

proton. Lewis further generalised the definition of an acid as an electron pair acceptor

and a base as an electron pair donor. The expressions for ionization (equilibrium)

constants of weak acids (Ka) and weak bases (Kb) are developed using Arrhenius definition.

The degree of ionization and its dependence on concentration and common ion are

discussed. The pH scale (pH = -log[H+]) for the hydrogen ion concentration (activity) has

been introduced and extended to other quantities (pOH = – log[OH–]) ; pKa = –log[Ka] ;

pKb = –log[Kb]; and pKw = –log[Kw] etc.). The ionization of water has been considered and

we note that the equation: pH + pOH = pKw is always satisfied. The salts of strong acid

and weak base, weak acid and strong base, and weak acid and weak base undergo

hydrolysis in aqueous solution.The definition of buffer solutions, and their importance

are discussed briefly. The solubility equilibrium of sparingly soluble salts is discussed

and the equilibrium constant is introduced as solubility product constant (Ksp). Its

relationship with solubility of the salt is established. The conditions of precipitation of

the salt from their solutions or their dissolution in water are worked out. The role of

common ion and the solubility of sparingly soluble salts is also discussed.