Equilibrium Class 11 Notes Chemistry Chapter 6 - CBSE

Chapter : 6

What Are Equilibrium ?

Chemical Equilibrium

Chemical Equilibrium is defined as the state at which there is no further change in concentration of reactants and products. At equilibrium the rate of forward reaction is equal to the rate of backward reaction.


Equilibrium Mixture

Equilibrium mixture is the mixture of reactants and products in the equilibrium state.

Based on the extent to which the reactions proceed to reach the state of chemical equilibrium, these may be classified in three groups.

  • The reactions which proceed almost to completion and the concentrations of the reactants left are negligible.
  • The reactions in which most of the reactants remains unchanged, i.e. only small amounts of products are
  • The reactions in which the concentrations of both the reactants and products are comparable when the system is in equilibrium.

Equilibrium In Physical Processes

Solid-Liquid Equilibrium

The equilibrium is represented as:


Rate of melting of ice = Rate of freezing of water.

The system here is in dynamic equilibriums and following can be inferred.

  • Both the opposing processes occur simultaneously.
  • Both the processes occur at the same rate so that the amount of ice and water – remains constant.

Liquid-Vapour Equilibrium

The equilibrium can be represented as:

Rate of evaporation = Rate of condensation When there is an equilibrium between liquid and vapours, it is called liquid-vapour equilibrium.

Solid-vapour Equilibrium

This type of equilibrium is attained where solids sublime to vapour phase. For example, when solid iodine is placed in a closed vessel, violet vapours start appearing in the vessel whose intensity increases with time and ultimately, it becomes constant. The equilibrium may be represented as:


Rate of sublimation of solid I2 to form vapour = Rate of condensation of I2 vapour to give solid I2

Equilibrium Involving Dissolution Of Solid In Liquid

  • Solution: When a limited amount of salt or sugar or any solute dissolves in a given amount of water solution
    is formed.
  • Saturated Solution: At a given temperature state is reached when no more solute can be dissolved then
    the solution is called saturated solution.

The equilibrium between a solid and its solution is indicated by the saturated solution and may be represented as:

Rate of dissolution = Rate of precipitation

$$\text{Sugar(s)}\rightleftharpoons\text{Sugar (in solution)}$$

Dissolution and precipitation takes place with the same speed. On adding a small amount of radioactive sugar to the saturated solution, the sugar present in the solution as well as in the solid state becomes radioactive.

Equilibrium Between A Gas And Its Solution In Liquid

The equilibrium arises due to difference in solubility of gases at different pressures. There is equilibrium between the molecules in the gaseous state and the molecules dissolved in the liquid under pressure. This equilibrium is governed by Henry’s law, which states that,

The mass of a gas dissolved in a given mass of a solvent at any temperature is proportional to the pressure of the gas above the solvent.”

Mathematically, m ∝ p

m = KHp (where KH = Henry’s Constant)

Some Features of Physical Equilibria

Process Conclusion
$$\text{Liquid}\rightleftharpoons\text{Vapour}\\\text{H}_2\text{O}\text{(l)}\rightleftharpoons \text{H}_2\text{O (g)}$$ PH2O constant at given temperature
$$\text{Solid}\rightleftharpoons\text{Liquid}\\\text{H}_{2}\text{O}\rightleftharpoons\text{H}_{2}\text{O}\space\text{(l)}$$ Melting point is fixed at constant pressure
$$\text{Solute}\rightleftharpoons\space\text{Solute (solution)}\\\text{Sugar}\rightleftharpoons\text{Sugar (solution)}$$ Concentration of solute in solution is constant at a given temperature
$$\text{Gas (g)}\rightleftharpoons\text{Gas (aq)}\\\text{CO}_{2}(g)\rightleftharpoons\space\text{CO}_2\text{(aq)}$$ [Gas (g)]/[Gas (aq)] is constant at a given temperature
[CO2 (aq)]/[CO2 (g)] is constant at a given temperature

Characteristics Of Equilibria Involving Physical Processes

  • The equilibrium can be attained only in closed systems at a given temperature.
  • At the equilibrium the measurable properties of the system remain constant.
  • The equilibrium is dynamic since both the forward and backward processes occur at same rate.
  • At equilibrium, the concentrations of substances become constant at constant temperature.
  • The value of equilibrium constant represents the extent to which the process proceeds before equilibrium is

Equilibrium In Chemical Processes

Equilibrium can also be achieved in chemical process involving reversible chemical reactions carried in closed container.

$$\textbf{A + B}\rightleftharpoons\textbf{C + D}$$

The dynamic nature of chemical equilibrium can be demonstrated in the synthesis of ammonia by Haber’s process. Haber started his experiment with the known amounts of N2 and H2 at high temperature and pressure. After a certain time he found that the composition of mixture remains the same even though some of the reactants are still present. This constancy indicates the attainment of equilibrium. In general, for a reversible reaction the chemical equilibria can be shown by :

for a reversible reaction the chemical equilibria can be shown by:

According to the equilibrium law,

$$\frac{[\text{C}][\text{D}]}{[\text{A}][\text{B}]} = \text{K}_{c}$$

Where, Kc is called the equilibrium constant.

For a general reaction,

$$\text{aA + bB}\rightleftharpoons\space\text{cC + dD.}$$

$$\text{K}_{c} =\frac{[\text{C}]^{c}[\text{D}]^{d}}{[\text{A}]^{a}[\text{B}]^{b}}$$

Equilibrium Constant for the reverse reaction is the inverse of the equilibrium constant for the reaction

$$\text{H}_{2} +\text{I}_{2}\rightleftharpoons\text{2HI}\space\text{is Kc,}\\\text{then}\space \text{2HI}\rightleftharpoons\space\text{H}_{2} +\text{I}_{2}\space\text{is 1/K}_{c}.$$

Law Of Chemical Equilibrium And Equilibrium Constant

At a given temperature, the product of concentrations of the reaction products raised to the respective stoichiometric coefficient in the balanced chemical equation divided by the product of concentrations of the reactants raised to their individual stoichiometric coefficients has a constant value. This is known as the Equilibrium Law or Law of Chemical Equilibrium.

The Norwegian chemists Cato Maximillian Guldberg and Peter Waage proposed in 1864 that the concentrations in an equilibrium mixture are related by the following equilibrium equation,

$$\text{K}_{c} =\frac{[\text{C}][\text{D}]}{[\text{A}][\text{B}]}$$

Where, Kc is the equilibrium constant

The equilibrium equation is also known as the law of mass action.

Relations between Equilibrium Constants for a General Reaction and its Multiples

Chemical Equation Equilibrium Constant
$$\text{aA + bB}\rightleftharpoons\text{c C + dD}$$ K
$$\text{cC + dD}\rightleftharpoons\text{aA +bB}$$ $$\text{K}'_c = \bigg(\frac{l}{\text{k}_{c}}\bigg)$$
$$\text{na A + nb B}\rightleftharpoons\text{nc C + nd D}$$ $$K^{''}_{c} = (K_{c}^{''})$$

Equilibrium In Homogeneous & Heterogeneous System

When in a system involving reversible reaction, reactants and products are in the same phase, then the system is called as homogeneous system.

Equilibrium in a system having more than one phase is called heterogeneous equilibrium.

Applications Of Equilibrium Constants

  • Expression for equilibrium constant is applicable only when concentrations of the reactants and products have attained constant value at equilibrium state.
  • The value of equilibrium constant is independent of initial concentrations of the reactants and products.
  • Equilibrium constant is temperature dependent, i.e., it has one unique value for a particular reaction
    represented by a balanced equation at a given temperature.
  • The equilibrium constant for the reverse reaction is equal to the inverse of the equilibrium constant for the forward reaction.
  • The equilibrium constant K for a reaction is related to the equilibrium constant of the corresponding reaction and its equation is obtained by multiplying or dividing the equation for the original reaction by a small integer.

Calculating Equilibrium Concentrations

  • Step 1: Write the balanced equation for the reaction.
  • Step 2: Make a table that lists the initial concentration, the change in concentration and the equilibrium
    concentration for each substance involved in the balanced reaction.
  • Step 3: Substitute the equilibrium concentrations into the equilibrium equation for the reaction and solve
    for x.
  • Step 4: Calculate the equilibrium concentrations from the calculated value of x.
  • Step 5: Check your results by substituting them into the equilibrium equation.

Relationship Between Equilibrium Constant K, Reaction Quotient Q And Gibbs Energy G

A mathematical expression of thermodynamic view of equilibrium can be described by tine equation.

∆G = ∆G0 + RT InQ

where, ∆G0 is standard Gibbs energy.

At equilibrium, ∆G = 0 and Q = Kc, so the equation becomes,

0 = ∆G0 + RT In Kc or, ∆G0 = -RT InKc

On changing the base, we get ∆G0 = -2.303RT logKc

We know that for a spontaneous process ∆G should be negative. So the value of Kc should be positive.

Le Chatelier’s Principle

Le Chatelier’s principle states that, “a change in any of the factors that determine the equilibrium conditions of a system will cause the system to change in such a manner so as to reduce or to counteract the effect of the change.” This is applicable to all physical and chemical equilibria.

Effect Of A Catalyst

Since catalyst increases the speed of both the forward and backward reactions to the same extent in a reversible reaction. Thus, catalyst has no effect on the equilibrium composition of a reaction mixture.

Arrhenius Concept Of Acids And Bases

  • Acids: According to Arrhenius theory, acids are substances that dissociates in water to give hydrogen ions H+ (aq).
  • Bases: Bases are substances that produce OH(aq) after dissociation in water.

$$\textbf{CH}_{3}\textbf{COOH} +\textbf{H}_{2}\textbf{O}\rightleftharpoons\space\\\textbf{CH}_{3}\space\textbf{COO}^{-} +\textbf{H}_{3}\textbf{O}^{\normalsize+}$$

Limitations of the Arrhenius Concept

  • According to the Arrhenius concept, an acid gives H+ ions in water but the H+ ions does not exist independently because of its very small size (~H-18 m radius) and intense electric field.
  • It does not account for the basicity of substances like, ammonia which does not possess a hydroxyl group.

The Bronsted-lowry Acids And Bases

According to Brönsted-Lowry theory, acid is a substance that is capable of donating a hydrogen ion H+ and bases are substances capable of accepting a hydrogen ion, H+ . The acid-base pair that differs only by one proton is called a conjugate acid-base pair. If Bronsted acid is a strong acid then its conjugate base is a weak base and vice versa.

Lewis Acids And Bases

G.N. Lewis in 1923 defined an acid as a species which accepts electron pair and base which donates an electron pair. Electron deficient species like AlCl3 , BH3 , H+ etc. can act as Lewis acids while species like H2 0, NH3 etc. can donate a pair of electrons, can act as Lewis bases.

Ionic Equilibrium In Solution

  • Electrolytes are substances which conduct electricity in their aqueous solution.
  • Strong Electrolytes are those electrolytes which on dissolution in water are ionized almost completely are
    called strong electrolytes.
  • Weak electrolyte are those electrolytes which on dissolution in water partially dissociated are called weak electrolyte.
  • Ionic Equilibrium is the equilibrium formed between ions and unionised substance. e.g.,

$$\textbf{Ice}\xrightarrow{}\textbf{Cubes water}$$

Acids, Bases And Salts

  • Acids are the substances which turn blue litmus paper to red and liberate dihydrogen on reacting with some
  • Bases are the substances which turn red litmus paper blue. It is bitter in taste. Common Example: NaOH,
    Na2 C03.
  • Salts are formed when acids and bases are mixed in the right proportion to react with each other. Some commonly known examples of salts are sodium chloride, barium sulphate, sodium nitrate.

The pH Scale

The pH of a solution is defined as the negative logarithm to base 10 of the activity (aH+) of hydrogen ion.

pH = -log aH+ = -log {[H+] / mol L-1}

Acidic solution has pH < 7, Basic solution has pH > 7, Neutral solution has pH = 7

Common Ion Effect

If in a aqueous solution of a weak electrolyte, a strong electrolyte is added having an ion common with the weak electrolyte, then the dissociation of the weak electrolyte is decreased or suppressed. The effect by which the dissociation of weak electrolyte is suppresed is known as common ion effect.

Di And Polybasic Acids

Acids which contain more than one ionizable proton per molecule are called Dibasic acids or polybasic acids or polyprotic acids. Common examples are oxalic acid, sulphuric acid, phosphoric acid etc.