Equilibrium Class 11 Notes CBSE Chemistry Chapter 7

Equilibrium Class 11 Notes for Chemistry Chapter 7 

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Access Class 11 Chemistry Chapter 7 – Chemical Equilibrium Notes

Introduction:

A state in chemical and biological reactions, during which there is no net change in the amounts of substances involved in these reactions is called the state of equilibrium. If water is boiled in a closed vessel, the rate of evaporation and rate of condensation will be equal in equilibrium conditions. The equation is as follows:

H2O (l) ⇌ H2O (vap)

Equilibrium is shown by a double arrow that is simultaneously moving in both directions. There is a significant amount of activity at the boundary between the liquid and vapour, thus this is not static equilibrium. As a result, at equilibrium, the rates of condensation and evaporation are equal. Equilibrium mixture refers to the combination of reactants and products in the equilibrium state.

Equilibrium in the Physical Process:

Transformations in nature may occur as physical changes involving the three states of matter; solid, liquid and gas. They may also undergo equilibrium when changing from one form to another.

1.Solid ⇌ Liquid  – Ice and water are in equilibrium with one another at a particular temperature and pressure. These temperatures are the normal melting or boiling points of a particular state of matter. 

A system is at equilibrium when both processes of melting and freezing occur simultaneously and at the same rate to keep the same amount of water and ice in reaction.

2. Liquid ⇌ Gas – Molecules from the vapour phase enter the liquid phase until an equilibrium is reached, in which case: 

Rate of Condensation = Rate of Evaporation,

This rate depends on the nature of the liquid, the amount of liquid and the applied temperature.

3.Solid ⇌ Gas – In sublimation reactions, solids turn into vapour and thereby reach a state of equilibrium. Eg. Camphor (solid) ⇌ Camphor (vapour)

General Characteristics of Equilibrium Involving Physical Processes:

• At a given temperature, equilibrium can be achieved only in closed systems.
• If antagonistic reactions occur at the same rate then the system is said to be dynamic. 
• Quantifiable properties in a system remain constant. 
• Equilibrium for physical processes can be determined by constant values of one of its parameters. 

Equilibrium in Chemical Processes:

Chemical equilibrium is reached when the rates of both forward and backward reactions are the same. 

Reversible Reactions

A given reaction, 

A + B ⇌  C + D

Is reversible in the sense that as the concentration of reactants decreases, the concentration of products formed will increase simultaneously. 

Irreversible Reactions

In the given reaction, 

A + B → C+ D

When the formed products cannot return to their reactant stages under the same conditions, it is an irreversible reaction. Such reactions do not show equilibrium and have a minimum Gibbs free energy.

Dynamic Nature of Chemical Equilibrium:

The reactants N2 and H2 in known quantities achieve equilibrium at a specific temperature during Haber’s process for producing ammonia. At equilibrium, N2, H2,  and NH3  are in stable concentrations.

Characteristics of Chemical Equilibrium:

• During equilibrium, reactants and products have constant concentrations.
• The dynamic nature of equilibrium is characterised by the equal rates of forwarding and backward reactions.
• Equilibrium can be achieved only if no solid products are permitted to get out of the reaction. 

Equilibrium Constant:

In a general reversible reaction, where A & B are reactants and C & D are products, 

A + B ⇌ C + D

Equilibrium constant,  Kc = [C] [D]

                                                 [A] [B]

Unit of Kc (Concentrations) – molL-1

The Law of Chemical equilibrium states that, at a given temperature, the product of concentrations of the reaction products raised to their respective stoichiometric coefficients in the balanced chemical equation divided by the product of concentrations of the reactants raised to their stoichiometric coefficients has a constant value. 

For a general reversible reaction, 

aA + bB ⇌ cC + dD 

Kc = [C]c[D]d / [A]a[B]b 

Where, A, B, C, and D are the equilibrium concentrations of the reactants and products

The product of the molar concentrations of the reactants, raised to their stoichiometric coefficients, divided by the product of the molar concentrations of the products, each raised to their stoichiometric coefficients, yields the equilibrium constant, which is a constant value.

Characteristics of Equilibrium Constant:

• The value of the equilibrium constant for a given reaction does not change even if the direction from which the equilibrium was approached is changed. It is independent of the concentration of reactants with which the calculation is started. It depends only upon temperature. 
• If a reaction is reversed, the value of KC is inverted.
• When the equation is divided by 2, Kc = √K
• When the equation is multiplied by 2, the new value of Kc = K2
• The inclusion of a catalyst will not affect the value of equilibrium. 

Predicting the Extent of Reaction:

The values of Kc help predict the dominant part of reactions :

1. If Kc > 103 has high values, the reaction is forward dominant
2. If Kc < 10-3 then the reaction is backwards dominant
3. Moderate values ( 10-3 > Kc > 103 ) signify equilibrium in neither direction. 

Predicting the Direction of the Reaction- Reaction Quotient:

The reaction quotient or QC helps to predict the direction of the reaction. It may not give us equilibrium values. 

aA + bB ⇌ cC + dD

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

Reaction proceeds in :

• The direction of reactants, when QC > KC
• The direction of products when QC < KC
• If QC = KC, the reaction mixture is already at equilibrium

Calculating Equilibrium Concentrations:

If initial concentrations are known, some steps can lead to finding equilibrium concentrations.

Step 1. A balanced equation is written

Step 2. Below the balanced equation, a list of each substance involved in that reaction is made

1. Initial concentration
2. Change in concentration on achieving equilibrium
3. The equilibrium concentration. 

Assign x as the concentration (mol/L) of one of the substances that react after reaching equilibrium, and then use stoichiometry of the reaction to determine the concentrations of the other substances to find x. 

Step 3. Replace the equilibrium concentrations in the reaction’s equilibrium equation and find x.

Step 4. Calculate the equilibrium concentrations from the calculated value of x.

Step 5: Use the equilibrium equation to substitute the findings to confirm.

Relationship Between Equilibrium Constant and Gibbs Free Energy:

A change in Gibbs’s free energy determines the spontaneity of a reaction. Kc does not depend on the rate of reaction. If ΔG denotes a change in Gibbs free energy, the following statements hold value. 

• If ΔG  is negative, the reaction is spontaneous in the forward direction
• If ΔG is positive, the reaction is spontaneous in the backward direction
• If ΔG is zero, the reaction is at equilibrium.

Mathematically,

ΔG = ΔG0 + RT ln Q

Where Q- reaction quotient and ΔG0 are the standard Gibbs free energy. 

At equilibrium, ΔG0 =0 and QC = KC

ΔG = ΔG0 + RT ln KC = 0

ΔG0 = – RT ln KC 

ln KC  = ΔG0 / – RT 

K = e-ΔG° / – RT 

Homogeneous Equilibria: 

A homogenous system is created when all reactants and products are in the same phase. 

N2 (g) + 3H2  ⇌ 2NH3 (g)    [all substances involved in the reaction are in gaseous phase] 

Heterogeneous Equilibria:

A heterogeneous system has many phases in the reaction.

H2O (l) ⇌ H2O (vap)            [This reaction takes place in a closed container]

Le Chatelier’s Principle:

According to the principle, if any of the variables that influence a system’s equilibrium circumstances to change, the system will adapt in a way that will lessen or cancel out the effects of the change.

Many factors influence this.

• Concentration change
• Temperature change
• Pressure change
• Volume change
• Involvement of a Catalyst
• Inert gas addition 

Ionic Equilibrium in Solution:

The establishment of equilibrium between ions in aqueous solutions is ionic equilibrium. 

A classic example of ionic equilibrium is that of NaCl dissociation to form Na+ and Cl– ions. Complete ionisation is observed in this case as NaCl is a strong electrolyte, unlike some ( Eg. Partial dissociation of CH3COOH, a weak electrolyte ) 

Michael Faraday classified substances into electrolytes and non-electrolytes.

Classification of Electrolytes:

Electrolytes can be classified into two types based on their strength. 

1. Strong Electrolytes – These undergo complete dissociation and give out ions. Eg. HCl, NaCl, NaOH
2. Weak electrolytes – These undergo partial dissociation into ions. Eg. CH3 COOH, NH4OH, HCN

 Acids, Bases and Salts:

This concept can be explained with multiple theories. Some of them are as follows.

• Arrhenius Concept of Acids and Bases:

Arrhenius claimed that in their aqueous solutions, bases produce hydroxyl ions while acids produce hydrogen ions. Acids produce H+ ions and bases produce OH– when present in an aqueous solution.

Arrhenius Acids:

Arrhenius suggested that acids are those which increase H+ concentration in water.

For an acid HX, 

HX (aq) → H+ (aq) + X– (aq), an acid molecule ionises in water.

Eg. HCl, CH3COOH

Arrhenius Bases:

Bases are those that increase OH– concentration in water, as suggested by Arrhenius. 

For a base MOH, 

MOH (aq) → M+ (aq) + OH– (aq), a base molecule ionises in aqueous solution. 

Eg. NaOH, NH4OH

• The Bronsted-lowry Acids and Bases:

According to Brönsted-Lowry, while a base is a proton acceptor, an acid is a proton giver. When a Brönsted-Lowry acid reacts with a base, it produces both the conjugate base for the reacting base and the conjugate acid for that base. A conjugate pair of acid and base therefore only varies by one proton.

The Bronsted-Lowry theory explains acids are substances that donate H+ ions and bases accept H+ ions. 

Dissociation of NH3 in water yields NH4 as the conjugate acid and OH– is the conjugate base. A water molecule that donates H+ is called a Lowry-Bronsted acid and the ammonia molecule that accepts H+ is the Lowry Bronsted base. 

• Lewis Acids and Bases:

Lewis expanded on the idea that a base is an electron pair donor and an acid is an electron pair receiver. This theory suggests that acids accept a pair of electrons and bases donate the same. 

BF3 (e– deficient ) + ：NH3 (e– donating) → BF3  ：+ NH3 

Examples: Lewis acids (AlCl3 , Co+3 , Mg+2 ) and Lewis bases ( H2O , NH3  , OH– )

The Ionisation Constant of Water and its Ionic Product:

Water molecules act as acids and bases. This can be represented as the following reaction and the dissociation constant K is given as :

H2O (l) (acid) + H2O (l) (base) ⇌ H3O+ (aq) + OH– (aq)

K =  [H3O+][ OH–] / [H2O]  ,where concentration of [H2O] remains constant since it is pure water.

In that case, this equation is known as the ionic product of water, Kw

KW = [H+][OH–]

Equal amounts of H+ and OH– are produced when water is dissociated. 

[H+]=[OH–]= 1 x 10-7 M

Kw = [H3O+][OH–] = (1 x 10-7)2 = 1 x 10-14 M2

Acidic – [H3O+] > [OH–]

Basic – [H3O+] < [OH–]

Neutral – [H3O+] = [OH–]

The pH Scale: 

The pH scale is a logarithmic scale representing the hydronium ion concentration in molarity

pH = – log [H+]

At 25℃, for pure water, [H+] = 10-7 M

Hence, pH (pure water) = – log (10-7) = 7  (Neutral solutions)

Acidic solutions generally possess an OH– ion concentration of more than 10-7 M and hence have a pH of less than 7. Basic solutions have an OH– ion concentration of less than 10-7 M, therefore their pH is greater than 7. 

pH scale goes up to 14 as the ionic product of water, Kw = [H3O+][OH–] = 10-14

Taking negative logarithms, 

-log Kw = -log {[H3O+][OH–]} = -log 10-14

pKw = pH + pOH = 14 

Ionisation Constants of Weak Acids:

For a weak acid HX, ionised partially in an aqueous solution, the reaction is as follows :

HX (aq) + H2O (l)  ⇌ H3O+ (aq) + X– (aq)

Equilibrium can be expressed in terms of the acid dissociation constant or ionisation constant (Ka)

Ka = C2? 2 / C (1-? ) = C? 2 / 1 – ? 

(where C – initial concentration of base,  ? = degree of ionisation of acid) 

Ionisation Constants of Weak Bases:

For a weak base MOH, the ionisation reaction is as follows :

MOH  ⇌ M+ (aq) + OH– (aq)

This shows a partial ionisation of MOH into M+ and OH–, like in an acid-dissociation equilibrium. For base ionisation, the equilibrium constant is called the base ionisation constant (Kb) which is expressed in terms of molality. 

Kb = [M]+ [OH–]  /   [MOH]

When equilibrium is reached, Kb can be written as :

Kb = C2 ?2 / C (1- ? ) = C?2  / 1- ?       

(where C – initial concentration of base,  ? = degree of ionisation of base)

pKb = -log (Kb)

Hydrolysis of Salts:

Reactions that form salts take place between acids and bases in redefined proportions. These salts get ionised in water, upon which the cations or anions formed exit as hydrated ions in aqueous solutions. Upon reaction with water, these salts may form corresponding acids or bases.

Eg. CH3COONa (Weak-acid and strong base), NH4Cl (Strong acid and weak base), CH3COONH4 (Weak acid and weak base)

Buffer Solutions:

Controlling pH is crucial in balancing chemical and biochemical processes. Particular pH values are to be maintained in many medical formulations. Buffer Solutions are solutions that resist change in pH when diluted or when modest volumes of acid or alkali are added.

Solubility Equilibria of Sparingly Soluble Salts: 

Salts have characteristic solubilities which depend on temperature and are classified as follows :

• If solubility > 0.1M then salt is soluble
• If 0.01 M < Solubility < 0.1M, salt is slightly soluble
• If solubility < 0.01 M, salt is sparingly soluble

Solubility product, Ksp = [Ag+] [Cl– ] = I. P = Ionic Product

If ionic product < Ksp, more salt can be dissolved through forwarding reaction

If ionic product > Ksp, precipitation of solid  salt takes place as a backward reaction

If ionic product = Ksp, no more salt can be dissolved, this is the saturation point.

Chemistry Class 11 Chapter 7 Notes 

Important concepts like reversible and irreversible reactions, chemical equilibrium, calculation of equilibrium constants, the rate of a reaction, Le Chatelier’s principle and electrolytes are explained in detail in this Extramarks’ Notes. Students can easily access these notes at their convenience and step up their exam preparations.  These solutions are provided here in accordance with the NCERT textbook and CBSE guidelines/ curriculum. These notes are prepared in such a way that they can get to the point answers without wasting much time on a single subject. Notes and revisions make them aware of their mistakes through guided practice and help to get the best results