Unit 7 – Equilibrium – Summary

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 ⇔ 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.

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