NCERT Notes For Class 11 Chemistry Chapter 7 Equilibrium

Class 11 Chemistry Chapter 7 Equilibrium

NCERT Notes For Class 11 Chemistry Chapter 7 Equilibrium, (Chemistry) exam are Students are taught thru NCERT books in some of state board and CBSE Schools.  As the chapter involves an end, there is an exercise provided to assist students prepare for evaluation.  Students need to clear up those exercises very well because the questions withinside the very last asked from those. 

Sometimes, students get stuck withinside the exercises and are not able to clear up all of the questions.  To assist students, solve all of the questions and maintain their studies without a doubt, we have provided step by step NCERT  Notes for the students for all classes.  These answers will similarly help students in scoring better marks with the assist of properly illustrated Notes as a way to similarly assist the students and answering the questions right.

NCERT Notes For Class 11 Chemistry Chapter 7 Equilibrium

Class 11 Chemistry Chapter 7 Equilibrium


Reversible and irreversible reactions

            A reaction that takes place in only one direction is called an irreversible reaction. e.g. Reaction between NaOH and HCl

           NaOH + HCl  →  NaCl + H2O

Reversible reactions are those which take place in both directions. i.e. here reactants combined to form products and the products recombine to form reactants.

E.g. Haber process for the preparation of ammonia   

The process by which reactants are converted to products is called forward reaction and the process by which products recombine to form reactants is called backward reaction. After sometimes, the rate of forward reaction becomes equal to the rate of backward reaction and the reaction attains equilibrium. Thus equilibrium is a state in which the rates of forward and backward reactions are equalEquilibrium is dynamic in nature. i.e. at equilibrium the reaction does not stop. The reactant molecules collide to form products and the product molecules collide to form the reactants and the rates of these reactions are equal.

Physical equilibrium: Equilibrium involving physical process is called physical equilibrium. E.g. melting of ice, evaporation of water, dissolution of solids or gases in liquids, sublimation etc.

            For example when ice and water are kept in a perfectly insulated thermoflask at 273K and 1 atm pressure, there exist an equilibrium between ice and water. At this stage, there is no change in the mass of ice and water. i.e. the rate of transfer of molecules from ice to water and the reverse process are equal. 

Chemical Equilibrium

Equilibrium associated with chemical reactions is called chemical equilibrium. At equilibrium, the concentrations of reactants and products are constant.

As the reaction proceeds, the concentration of the reactants decreases and that of the products increases. After sometimes, the two reactions occur at the same rates and an equilibrium state is reached. This can be illustrated by the following graph.

            After the equilibrium is attained, the concentration of the reactants and products become constant. At equilibrium, the rate of forward reaction (rf) = the rate of backward reaction (rb).

Law of Chemical Equilibrium and Equilibrium Constant

            This law was proposed by Guldberg and Waage. It states that at constant temperature, the product of concentration of the products to that of the reactants, in which each concentration terms is raised to a power which is equal to the stoichiometric coefficients in the balanced chemical equation, has a constant value.

             For a general reversible reaction:

Equilibrium constant for gaseous reactions

Where Kp is called equilibrium constant in terms of partial pressure, PHI, PH2 and PI2 are the partial pressures of HI, H2 and I2 respectively.

Relation between Kc and Kp

From ideal gas equation, PV = nRT, P = nRT/V = CRT (since n/V = C, the concentration)

Therefore, PA = CART, PB = CBRT, PC = CCRT and PD= CDRT Substitute these values in equation (2), we get

And ∆n is the change in no. of moles of gaseous species. i.e. ∆n = no. of moles of gaseous products – no. of moles of gaseous reactants.

Special cases:

  1. If ∆n = 0, then Kp = Kc
  2. If ∆n > 0, then Kp > Kc and 
  3. If ∆n < 0, then Kp < Kc

Homogeneous and heterogeneous equilibria

            An equilibrium reaction in which all the reactants and products are in the same phase is called homogeneous equilibrium.

Equilibrium reaction in which the reactants and products are in different phases is called heterogeneous equilibrium.

Characteristics of Equilibrium constant

The important characteristics of equilibrium constant are:

  1. Equilibrium constant is applicable only when the concentrations of the reactants and products have attained their equilibrium state.
  2. The value of equilibrium constant is independent of the initial concentrations of reactants and products.
  3. The value of equilibrium constant depends on temperature.
  4. The equilibrium constant for the reverse reaction is the reciprocal of that of the forward reaction.

Applications of equilibrium constant

             The important applications of equilibrium constant are:

  • Prediction of the extent of a reaction

Greater the value of equilibrium constant, greater will be the concentration of products. In general, 

  1. If Kc > 103 (i.e. Kc is very large), the reaction proceeds nearly to completion
  2. If Kc < 10-3 (i.e. if Kc is very small), the reaction proceeds rarely.
  3. If the value of Kc is in between 103 and 10-3 appreciable concentrations of both reactants and products are present.
  • Prediction of the direction of the reaction

By knowing the values of Kc and Qc, we can predict the direction of a reaction. The   reaction  quotient(Qc) is defined in the same way as the equilibrium constant (Kc) except that the concentrations in Qc are not necessarily the equilibrium values. 

If Qc > Kc, the reaction will proceed in the direction of reactants (reverse direction). If Qc < Kc, the reaction will proceed in the direction of products (forward direction). If Qc = Kc, the reaction mixture is at equilibrium.

  • Calculation of equilibrium concentrations

By knowing the value of equilibrium constant, we can calculate the equilibrium concentrations of  reactants and products.


          The Gibb’s energy change of a reaction is related to the reaction quotient (Q) by the equation:

                      ΔG = ΔG0 + RT lnQ

where, ΔG0 is standard Gibbs energy.

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

     0 = ΔG0 + RT ln Kc

Or,      ΔG0 = – RT lnKc 

On changing the base, we get ΔG0 = – 2.303RT logKc

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

Factors affecting equilibrium

  The important factors affecting equilibrium are temperature, pressure, concentration and catalyst. The effect of these factors on equilibrium state can be explained by using Le Chatlier’s Principle.  It states that whenever there is a change in concentration, pressure or temperature of a system at equilibrium, the system will try to readjust in such a way so as to cancel the effect of that change.

1. Effect of concentration change

  • If we change the concentration of reactants or products in an equilibrium process, then according  to Le Chatlier’s principle, the system will try to reduce the effect of that change.
  • For this the rate of either forward or backward reaction changes.
    • In general, an increase in concentration of reactants increases the rate of forward reaction (i.e. the equilibrium is shifted to the forward direction) and an increase in concentration of products increase the rate of backward reaction rate (i.e. the equilibrium is shifted to the backward direction).
  • If we remove SO3 from the reaction mixture, its concentration decreases. Here also to increase the concentration, the system will shift to the forward direction.
  • In Haber process for the preparation of ammonia, the amount of ammonia formed can be increased by increasing the concentration of N2 or H2 or by removing NH3 from the reaction mixture.

2. Effect of temperature change

  • According to Le Chatlier’s principle, increase in temperature favours endothermic process and decrease in temperature favours exothermic process.
  • In a reversible reaction, if the forward reaction is endothermic, the backward reaction will be exothermic.
  • Here since ∆H is negative, the forward reaction is exothermic. So to increase the production of NH3, temperature should decrease.
  • At higher temperature, the rate of backward reaction increases. i.e. the ammonia formed is decomposed to N2 and H2.

3. Effect of pressure change

  • Pressure has its role only in gaseous reactions. In general an increase in pressure favours the  reaction in which number of moles decreases and vice versa.
  • This can be achieved by shifting the reaction into the direction in which no. of moles decreases (since pressure is directly proportional to the number of moles).
  • So in this reaction, the equilibrium will shift to the forward direction. 
    • If the volume of the reaction mixture is halved, the concentration and the partial pressure become doubled.
  • So the reaction is shifted to the direction in which the number of moles or volume decreases.

4. Effect of catalyst

  • In an equilibrium reaction, a catalyst increases the rate of both forward and backward reactions  Simultaneously and helps to attain the equilibrium faster.
  • It lowers the activation energy for the forward and backward reactions by exactly the same amount.
  • It does not affect the equilibrium composition of the reaction mixture.
  • Thus in Haber process for the manufacture of ammonia, iron (Fe) acts as the catalyst.
  • In contact process for the manufacture of sulphuric acid, V2O5 (Vanadium pentoxide) is used as the catalyst for the conversion of SO2 to SO3.

5. Effect of inert gas addition

  • If an inert gas (which does not take part in the reaction) is added to an equilibrium mixture at constant volume (or at constant pressure), there is no change to the equilibrium.
  • It is because the addition of inert gas at constant volume (or at constant pressure) does not change the partial pressure or the concentration of the reactants and the products.

Ionic equilibrium in solution

Equilibrium involving ions is called ionic equilibrium. E.g. dissociation of acetic acid in water.

Electrolytes and non-electrolytes

Based on the ability to conduct electricity, Michael Faraday divided substances into two – electrolytes and non-electrolytes.

  • Electrolytes are substances which conduct electricity in molten state or in solution state. e.g. All acids, bases and almost all salts
  • Non- electrolytes are substances which do not conduct electricity in molten state or in solution state. e.g. sugar, urea etc.

Electrolytes are further classified into two – strong electrolytes and weak electrolytes. 

  • Strong electrolytes are electrolytes which dissociate almost completely in aqueous solution. 
  • E.g. strong acids like HCl, HNO3, H2SO4 etc., strong bases like NaOH, KOH etc. and salts like NaCl, KCl, Na2SO4, K2SO4, KNO3, NaNO3 etc.
  • Electrolytes which dissociate only partially in aqueous solution are called weak electrolytes. E.g. weak acids like CH3COOH, formic acid (HCOOH) etc., weak bases like Mg(OH)2, Ca(OH)2, NH4OH etc. and some salts like CaSO4, BaSO4 etc.
  • A weak electrolyte dissociates only partially in aqueous solution and so an equilibrium is formed between the ions and the unionised molecules.
  • This type of equilibrium involving ions in aqueous solution is called ionic equilibrium.

Acids, Bases and Salts

Acid – base concepts:

1.Arrhenius concept:

  • According to this concept acids are substances which give hydrogen ion (H+) or hydronium ion (H3O+) in aqueous solution and bases are substances which give hydroxyl ion (OH) in aqueous solution.

Limitations: this concept is applicable only to aqueous solutions. Also it could not account for the basicity of substances like NH3 which do not possess OH ion.

  • The Bronsted – Lowry concept:

According to this concept acids are proton (H+) donors and bases are (H+) acceptors.

  • Here NH3 is a base since it accepts an H+ ion to form NH4+ and H2O is an acid since it donates an H+ ion to form OH . In the reverse reaction, NH4+ is an acid and OHis a base.
  • The acid base pair that differs by only one proton is called a conjugate acid – base pair.
  • An acid formed from a base is called  conjugate acid and a base formed from an acid is called conjugate base.
  • If the acid is strong, its conjugate base is weak and vice versa. So in the above example Cl- is a weak conjugate base of the strong acid HCl. 

Water can act both as acid and base. So it is an amphoteric substance.  e.g.

3. Lewis concept:

  • According to this concept acids are electron pair acceptors and bases are electron pair donors. 
    • Substances which donate electron pair are called Lewis bases and substances which accept electron pair are called Lewis acids.
Example for Lewis acids are BF3, AlCl3, H+, Co3+, Mg2+ etc.   Example for Lewis bases are NH3, H2O, OH, Cl, Bretc.
  • For a substance to act as Lewis acid, it should contain vacant orbitals and for a substance to act as Lewis base, it should contain lone pairs of electrons.

            e.g. for acid – base reaction is      BF3 + NH3 → BF3←NH3

All cations are Lewis acids and all anions are Lewis bases.

The ionization constant of water (The ionic product of water)

Water is a weak electrolyte and hence it ionizes only partially as:

  • Where Kw is called ionization constant of water or ionic product of water.
  • It is defined as the product of the molar concentration of hydrogen ion (hydronium ion) and hydroxyl ion in water or in any aqueous solution.

          For pure water at 298K, [H+] = [OH] = 10-7M.

Therefore, Kw = [H+][OH] = 10-7 x 10-7 = 10-14M2

  • The value of Kw is temperature dependent. 

By knowing the concentrations of H3O+ and OH ions, we can predict the nature of an aqueous solution.

If [H3O+] > [OH], the solution is acidic

If [H3O+] < [OH], the solution is basic

If [H3O+] = [OH], the solution is neutral

The pH scale  

  • pH is defined as the negative logarithm of the hydrogen ion or hydronium ion concentration in moles per litre (i.e. molarity). i.e. pH = – log[H+] or pH = – log[H3O+] Negative logarithm of hydroxyl ion concentration in mol/L is called pOH. i.e. pOH = -log[OH]
  • For pure water, at 298K (250c), [H+] = 10-7. Therefore pH of pure water is 7.
  • The pH scale was introduced by Sorensen. It contains numbers from 0 to 14. If the pH is less than 7, the solution is acidic, if it is greater than 7, it is basic and if it is 7, the solution is neutral.
  • If the pH is 0, 1, 2 etc., the solution is a strongly acidic and if it is 12, 13, 14 etc, it is strongly basic.

The pH of our blood is 7.4 and that of our saliva is 6.4. So blood is slightly basic and saliva is slightly acidic.

Relation between pH and pOH

            We know that Kw = [H+][OH] = 10-14 at 298K.

Taking negative logarithm on both sides:

          -logKw = -log[H+] + -log[OH] = -log 10-14

Or,      pKw = pH + pOH = 14

Thus by knowing the pH, we can calculate pOH as pOH = 14 – pH

Ionisation constant of weak acids

Where Ka is the dissociation constant of the weak acid.

Larger the value of Ka, stronger is the acid. The negative logarithm of Ka is called pKa.  i.e. pKa = -logKa.

Greater the value of Ka, smaller will be the value of pKa and stronger will be the acid.

Ionisation constant of weak bases

As the value of Kb increases, pKb decreases and the basic strength increases.

Relation between Ka, Kb and Kw

            For conjugate acid – base pair Ka and Kb are related as:   Ka x Kb = Kw This can be deduced as follows:

           or,  pKa + pKb = pKw = 14           

Factors Affecting Acid Strength

  • The extent of dissociation of an acid and its acidic strength depend on the strength and polarity of the H-A bond.
  • Weaker the H-A bond, more easily it dissociates to give H+ ion and hence stronger is the acid.
  • Also, greater the polarity of the H – A bond (when the electronegativity difference between the atoms H and A increases) the more easily it dissociates and hence greater is the acidity.
  • In a group, acidity is mainly determined by bond strength.
  • From top to bottom in a group, as the size of A increases, H-A bond strength decreases and so the acid strength increases.
  • Thus acidic strength of hydrohalic acids increases in the order: HF < HCl < HBr < HI.
  • In a period, acidity is mainly determined by polarity of the bond.
  • As the electronegativity of A increases from left to right in a period, the strength of the acid also increases.
  • So the acidity of hydrides of second period elements increases in the order: CH4 < NH3 < H2O < HF.

Hydrolysis of Salts and the pH of their Solutions

  • The interaction of anion or cation or both of a salt with water is known as salt hydrolysis.
  • The cations (e.g., Na+, K+, Ca2+, Ba2+, etc.) of strong bases and anions (e.g., Cl, Br, NO3 , ClO4 etc.) of strong acids  do not get hydrolyse.
  • So the solutions of salts formed from strong acids and bases (e.g. NaCl, KCl, NaNO3, KNO3, Na2SO4, K2SO4 etc) are neutral i.e., their pH is 7.
  1. Hydrolysis of salt of strong base and weak acid:
    1. Sodium acetate (CH3COONa), sodium carbonate (Na2CO3), potassium cyanide (KCN) etc. are examples for such type of salts.
    1. Here only the anion of the weak base undergoes hydrolysis (since cation of the strong base does not hydrolyse).
    1. So the solution of such salts will be basic. i.e. pH > 7.
    1. pH of such salt solution is given by pH = 7 + ½ (pKa + log C) where C is the concentration of salt.
  2. Hydrolysis of salt of weak base and strong acid:
    • NH4Cl, NH4NO3, CuSO4 etc are examples for such type of solutions.
    • Here only cation of weak base undergoes hydrolysis.
    • So the solution is acidic. pH of such a solution is given by pH = 7 + ½ (pKb + log C).
  3. Hydrolysis of salt of weak base and weak acid:
    • ammonium acetate (CH3COONH4), ammonium carbonate [(NH4)2CO3] etc. are examples for such type of salts.
    • Here both cation and anion undergo hydrolysis and hence weak acid and weak base are produced in solution.
    • So the solution may be neutral, acidic or basic depending upon the relative strength of acid and base formed.
    • pH of such a solution is given by pH = 7 + ½ (pKa + pKb).

Common Ion Effect

  • It is the suppression of the dissociation of a weak electrolyte by the addition of a strong electrolyte containing a common ion. For e.g. consider the dissociation of acetic acid (a weak electrolyte).
  • If we add some sodium acetate (CH3COONa) to the above equilibrium reaction, the concentration of acetate ion increases.
  • Then according to Le-Chatlier’s principle, the equilibrium will shift towards left or, the rate of forward reaction decreases. i.e. the dissociation rate of acetic acid decreases.
  • This is known as common ion effect.

          Another example is the dissociation of the weak base ammonium hydroxide (NH4OH)

If we add some NH4Cl to the above equilibrium process, the concentration of NH4+ increases and hence the equilibrium shifted to the left. i.e. the dissociation rate of NH4OH decreases.

Buffer Solutions

  • Solution which resists the change in pH on dilution or with the addition of small amount of acid or alkali is called Buffer solution.
  • There are two types of buffer solutions – acidic buffer and basic buffer.
  • Acidic buffer is a mixture of a weak acid and its salt with a strong base. E.g. a mixture of acetic acid and sodium acetate acts as an acidic buffer around pH 4.75.
  • Basic buffer is a mixture of a weak base and its salt with a strong acid. E.g. a mixture of NH4OH and NH4Cl acts as a basic buffer around pH 9.25.

PH of a Buffer solution – Henderson – Hasselbalch Equation

Consider an acidic buffer prepared by mixing a weak acid HA and its conjugate bas A. The weak acid HA ionises in water as:

This equation is known as Henderson-Hasselbalch equation. The equation can be written in general form

Solubility Equilibrium

  • The solubility of a salt in solvent mainly depends on the lattice enthalpy and solvation enthalpy.
  • As a general rule, for a salt to be able to dissolve in a particular solvent its solvation enthalpy must be greater than its lattice enthalpy.
  • Each salt has its characteristic solubility which depends on temperature.

We can classify salts on the basis of their solubility in the following three categories.

Category ISolubleSolubility > 0.1M
Category IISlightly SolubleSolubility in between 0.01M and 0.1M
Category IIISparingly SolubleSolubility < 0.01M

Solubility Product Constant

Consider the dissociation of a sparingly soluble salt in water. Since it dissolves only partially, there exists an equilibrium between the undissolved solid and the ions. e.g. Solution of barium sulphate (BaSO4)

For a pure solid, the concentration remains constant.

Therefore, K.[BaSO4] = [Ba2+][SO42-]  Or, Ksp = [Ba2+][SO42-]

            Where Ksp is called the solubility product constant or simply the solubility product. It is defined as the product of the molar concentration of ions of a sparingly soluble salt or in a saturated solution. If S is the molar solubility of BaSO4, then  Ksp = S x S = S2

For a general salt AxBy, its dissociation can be denoted as:                

If the concentration in the above equation is not the equilibrium concentration, then Ksp is given by Qsp.

At equilibrium, Ksp = Qsp.

If Ksp > Qsp, the dissolution process occurs and if Ksp < Qsp, the precipitation of the salt occurs.

Solubility product and common ion effect found application in the purification of NaCl.

 If we take a saturated solution of NaCl and pass HCl gas through it, then NaCl gets precipitated due to the increased concentration of Cl ions. NaCl thus obtained has very high purity.

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