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Class 11 Chemical Equilibrium Important Short and Long Questions with Answers

The aqueous solution of sugar does not conduct electricity. However, when sodium chloride is added to water, it conducts electricity. How will you explain this statement on the basis of ionisation and how is it affected by the concentration of sodium chloride?

Answer:

Sugar is a non-electrolyte, and when it dissolves in water, there will be no ionisation. If there are no free ions, it will not conduct electricity. In contrast, When sodium chloride is added to water, it is wholly ionised to give Na⁺ ions and Cl⁻ ions. Due to the presence of ions, there will be electrical conductance as ions are carriers of electric current.

BF₃ does not have protons but still acts as an acid and reacts with NH₃. Why is it so? What type of bond is formed between the two?

Answer:

BF₃ is electron-deficient and hence acts as Lewis acid. : NH₃ has one lone pair, which can donate to BF₃ and form a coordinate bond with NH₃, i.e. (NH₃: → BF₃).

Conjugate acid of a weak base is always stronger. The decreasing order of basic strength of the following conjugate bases will be? OH^–, RO^–, CH₃COO^–, CI^–

Answer:

Conjugate acids of given bases are H₂O, ROH, CH₃COOH, and HCl.

Their acidic strength is in the order.

HCl > CH₃COOH > H₂O >ROH.

Hence, their conjugate bases will have strength in the order.

RO^– > OH^– > CH₃COO^– > CI^–

Arrange the following in increasing order of pH.

KNO₃ (aq), CH₃COONa (aq), NH₄Cl (aq), C₆H₅COONH₄ (aq)

Answer:

NH₄Cl < C₆H₅COONH₄ < KNO₃ < CH₃COONa

KNO₃ is formed from a strong acid, nitric acid and strong base, potassium hydroxide. Salts of strong acid and strong base hydrolyse to form a neutral solution. Thus its pH will be nearly 7. Sodium acetate is formed from a weak acid, acetic acid, and strong base sodium hydroxide. Thus its pH will be more than 7.C₆H₅COONH₄ is formed from a mild acid, C₆H₅COOH and a weak base, NH₄OH. Thus its pH will be less than 7. NH₄Cl is formed from a weak base, NH₄OH, and strong acid, HCl. Thus its pH value would be less than 5.

The value of Kc for the reaction 2HI (g) ⇋ H₂ (g) + I₂ (g) is 1 x 10^-4. At a given time, the composition of reaction mixture is [HI] = 2 x 10^-5 mol, [I₂ ] = 1 x 10^-5 mol. Determine the direction in which the reaction will proceed?

Answer:

At a given time, the reaction quotient Q for the reaction will be given by the expression.

Q = [H₂ ][I₂ ] / [HI]

Q = 1 X 10^-5 X 1 X 10^-5 / (2 X 10^-5 )²

Q = 1 / 4

Q = 0.25

Q = 2.5 X 10^-1

As the value of the reaction quotient is greater than the value of Kc , i.e. 1×10⁻⁴, the reaction will proceed in the reverse direction.

On the basis of the equation pH = – log [H⁺], the pH of 10^-8 mol dm^-3 solution of HCI should be 8. However, it is observed to be less than 7.0. Explain the reason.

Answer:

Here, the water concentration cannot be neglected since the solution is very dilute.

[H₃O⁺] = 10⁻⁸ +10⁻⁷ M

[H₃O⁺] = 10⁻⁸ (1 + 10)

[H₃O⁺] = 11 X 10⁻⁸ M

pH = – log [H₃O⁺]

pH = – log 11 X 10⁻⁸ M

pH = 8 – log 11

pH = 8 – 1.04

pH = 6.96

pH will be less than 7.0.

The pH of a solution of a strong acid is 5.0. What will be the pH of the solution obtained after diluting the given solution 100 times?

Answer:

pH = 5 means [H⁺] = 10⁻⁵

On diluting 100 times ,

[H⁺] = 10⁻⁵ / 100 = 10⁻⁷

On calculating the pH using the equation pH = − log [H⁺], the pH value comes out to be 7, which is impossible.

Hence, Total H⁺ ion concentration = H⁺ ions from acid + H⁺ ion from water

[H⁺] = 10⁻⁷ +10⁻⁷ M

[H⁺] = 2 X 10⁻⁷

pH = 7 – 0.3010

pH = 6.699

The pH of 0.08 mol dm^-3 HOCI solution is 2.85. Calculate its ionisation constant.

Answer:

pH of HOCl = 2.85

But, – pH = log [H⁺]

∴ -2.85 = log [H⁺]

-3.15 = log [H⁺] [H⁺] = 1.413 × 10^–3

For the weak monobasic acid [H⁺] = (K_a X C) ^{1 / 2}

K_a = [H⁺] ² / C

K_a = (1.413 × 10^–3 )² / 0.08

K_a = 24.957 × 10^–6

K_a = 2.4957 × 10^–5

Calculate the pH of a solution formed by mixing equal volumes of two solutions A and B of a strong acid having pH = 6 and pH = 4 respectively.

Answer:

The pH of solution A = 6,

Therefore, the concentration of [H⁺] ion in solution A=10⁻⁶ mol / L

The pH of solution B = 4

Therefore, the concentration of [H⁺] ion in solution B=10⁻⁴mol / L.

On mixing one litre of each solution, total volume = 1L + 1L = 2L.

Amount of H⁺ ions in 1L of solution A = concentration × Volume (V) =

Amount of H⁺ ions in 1L of solution A = 10⁻⁶ X 1 = 10⁻⁶

Amount of H⁺ ions in 1L of solution B =10⁻⁴ X 1 = 10⁻⁴

∴ The total amount of H⁺ ions in the solution formed by mixing solutions A and B is (10⁻⁵ mol + 10⁻⁴ mol)

This amount is present in 2L solution,

∴ Total [H⁺] = 10⁻⁴ (1 + 1.01) / 2 mol / L

Total [H⁺] = 1.01 X 10⁻⁴ / 2 mol / L

Total [H⁺] = 0.5 X 10⁻⁴ mol / L

Total [H⁺] = 5 X 10⁻⁵ mol / L

pH = – log [H⁺]

pH = – log [5 X 10⁻⁵]

pH = – [ log 5 – 5 log 10]

pH = – log 5 + 5

pH = 5 – log 5

pH = 5 – 0.6990

pH = 4.3010

pH = 4.3

Thus, the pH will be 4.3.

The solubility product of Al(OH)₃ is 2.7 x 10^-11. Calculate its solubility in g / L and also find out the pH of this solution. (Atomic mass of Al = 27 u).

Answer:

Let S be the solubility of Al(OH)₃

K_sp = [Al³⁺] [OH⁻]³

K_sp = (S) (3S)³

K_sp = 27S⁴

S⁴ = K_sp / 27

S⁴ = 27 × 10^-11 / 27 x 10

S⁴ = 1 x 10 ^-12 mol / L.

S = 1 x 10 ^-3 mol / L.

(i) Solubility of Al(OH)₃: Molar mass of Al(OH)₃ is 78g.

Therefore, the solubility of Al(OH)₃ in g / L = 1×10 ³ × 78 g / L

Solubility of Al(OH)₃ in g / L = 78 × 10⁻³ g / L

Solubility of Al(OH)₃ in g / L = 7.8 × 10⁻² g / L

(ii) pH of the solution: S = 1 × 10⁻³ mol / L

[OH] = 3S = 3 × 1 × 10⁻³ = 3 × 10⁻³

pOH =3 − log 3

pH = 14 − pOH = 11 + log 3 = 11.4771.

A reaction between ammonia and boron trifluoride is given below:

: NH₃ + BF₃ → H₃N: BF₃

Identify the acid and base in this reaction. Which theory explains it? What is the hybridisation of B and N in the reactants?

Answer:

NH₃ is Lewis base, while BF₃ is Lewis acid. Lewis’s electronic theory of acids and bases explains it.

The hybridisation state of nitrogen in NH₃ is sp³ hybridised, and boron in BF₃ is sp² hybridised.

Following data is given for the reaction: CaCO₃(s) → CaO (s) + CO₂ (g)

Δf H^– [CaO (s)] = – 635.1 kJ / mol

Δf H^– [CO₂ (g)] = – 393.5 kJ / mol

Δf H^– [CaCO₃ (s)] = – 1206.9 kJ / mol

Predict the effect of temperature on the equilibrium constant of the above reaction.

Answer:

Δf H^– = Δf H^– [CaO (s)] + Δf H^– [CO₂ (g)] − Δf H^– [CaCO₃ (s)]

∴ Δf H^– = 178.3 kJ / mol.

The reaction is endothermic.

Hence, according to Le Chatelier’s principle, the reaction will proceed in the forward direction on increasing temperature.

How can you predict the following stages of a reaction by comparing the value of Kc and Qc?

(i) Net reaction proceeds in the forward direction.

(li) Net reaction proceeds in the backward direction.

(iii) No net reaction occurs.

Answer:

The values of Kc and Qc are self-explanatory and less than or greater than one another decides the direction in which reaction will proceed as follows-

(i) As Qc < Kc, the reaction will proceed in the direction of products (forward direction).

(ii) If Qc > Kc, the reaction will proceed in the direction of reactants (reverse direction).

(iii) If Qc = Kc, no net reaction will occur.

On the basis of the Le Chatelier principle explain how temperature and pressure can be adjusted to increase the yield of ammonia in the following reaction.

N₂ (g) + 3H₂ (g) ⇋ 2NH₃ (g) ΔH = – 92.38 kJ mol^-1

What will be the effect of addition of argon to the above reaction mixture at constant volume?

Answer:

According to Le Chatelier’s principle, raising the temperature shifts the equilibrium to the left and decreases the equilibrium concentration of ammonia since it is an exothermic reaction. In other words, low temperature and high pressure are favourable for a high yield of ammonia. There will be no change in equilibria on the addition of argon (Ar).

A sparingly soluble salt having general formula A_x B_y and molar solubility S is in equilibrium with its saturated solution. Derive a relationship between the solubility and solubility product for such salt.

Answer:

A_x^p+ + B_y^q− ⇋ xA^p+ (aq) + yB^q− (aq)

S moles of A B dissolve to give x S moles of A^p+ and y S moles of Bq⁻.

K_sp = [A^p+] ^x [B^q−] ^y

K_sp =[x5] ^x [y5] ^y

x ^x y ^y 5 ^x+y.

Write a relation between ΔG and Q and define the meaning of each term and answer the following:

(a) Why a reaction proceeds forward when Q < K and no net reaction occurs when Q = K.

(b) Explain the effect of an increase in pressure in terms of reaction quotient O. for the reaction: CO (g) + 3H₂ (g) ⇋ CH₄ (g) + H₂O (g)

Answer:

(a) ΔG = ΔG^⊝ + RT ln Q

ΔG^⊝ = Change in free energy as the reaction proceeds,

ΔG = Standard free energy change,

Q = Reaction quotient,

R = Gas constant,

T = Absolute temperature.

Since ΔG^⊝ = − RTlnK

∴ ΔG = −RTlnK + RTlnQ

ΔG = RT lnQ / K

If Q < K, ΔG will be negative. The reaction will proceed in forward direction.

If Q = K, ΔG = 0, the reaction is in equilibrium, and no net reaction is there.

(b) CO (g) + 3H₂ (g) ⇋ CH₄ (g) + H₂O (g),

According to the Le Chatiliers principle, when the pressure increases, the equilibrium will shift in that direction where there are fewer gas molecules.

As per the available information,

The number of moles of reactants = 1 + 3 = 4

The number of moles of product = 1 + 1 = 2

Hence, the equilibrium will shift forward as the number of moles of product is less.

Important Topics Covered in Class 11 Chemistry of Chapter Equilibrium are listed below:

1. Solid-liquid Equilibrium
   • Liquid-vapour Equilibrium
   • Solid-vapour Equilibrium
   • Equilibrium Involving Dissolution of Solid or Gases in Liquids
   • General Characteristics of Equilibria Involving Physical Processes
2. Equilibrium in Chemical Processes – Dynamic Equilibrium
3. Law of Chemical Equilibrium and Equilibrium Constant
4. Homogeneous Equilibria
   • Equilibrium Constant in Gaseous Systems
5. Heterogeneous Equilibria
6. Applications of Equilibrium Constants
   • Predicting the Extent of a Reaction
   • Predicting the Direction of the Reaction
   • Calculating Equilibrium Concentrations
7. Relationship between Equilibrium Constant K, Reaction Quotient Q and Gibbs Energy G
8. Factors Affecting Equilibria
   • Effect of Concentration Change
   • Effect of Pressure Change
   • Effect of Inert Gas Addition
   • Effect of Temperature Change
   • Effect of a Catalyst
9. Ionic Equilibrium in Solution
10. Acids, Bases and Salts
    • Arrhenius Concept of Acids and Bases
    • The Bronsted-lowry Acids and Bases
    • Lewis Acids and Bases
11. Ionisation of Acids and Bases
    • The Ionisation Constant of Water and Its Ionic Product
    • The pH Scale
    • Ionisation Constants of Weak Acids
    • Ionisation of Weak Bases
    • The Relation Between K_a and K_b
    • Di- and Polybasic Acids and Di- and Polyacidic Bases
    • Factors Affecting Acid Strength
    • Common Ion Effect in the Ionisation of Acids and Bases
    • Hydrolysis of Salts and the Ph of Their Solutions
12. Buffer Solutions
13. Solubility Equilibria of Sparingly Soluble Salts
    • Solubility Product Constant
    • Common Ion Effect on Solubility of Ionic Salts

Brief Summary of Chapter Chemical Equilibrium

What is Chemical Equilibrium?

Chemical equilibrium is a state in a chemical reaction in which the concentrations of reactants and products remain constant over time. This does not mean that the forward and backward reactions have stopped, but rather that they are occurring at the same rate.

Types of Chemical Equilibrium

There are two main types of chemical equilibrium:

• Homogeneous equilibrium: This occurs when all of the reactants and products are in the same phase, such as a gas or a solution.
• Heterogeneous equilibrium: This occurs when the reactants and products are in different phases, such as a solid and a gas.

Characteristics of Chemical Equilibrium

Chemical equilibrium is characterized by the following:

• The concentrations of all reactants and products remain constant over time.
• The rate of the forward reaction is equal to the rate of the backward reaction.
• The system is dynamic, meaning that the forward and backward reactions are still occurring, but at the same rate.
• The system is reversible, meaning that the reaction can go in both the forward and backward directions.

Factors Affecting Chemical Equilibrium

The following factors can affect the position of equilibrium in a chemical reaction:

• Concentration: Changes in the concentration of reactants or products can shift the equilibrium position. For example, increasing the concentration of reactants will shift the equilibrium position to the right, favoring the formation of products.
• Temperature: Increasing the temperature of a chemical reaction will generally shift the equilibrium position to the right, favoring the formation of products that absorb heat.
• Pressure: Changing the pressure of a chemical reaction can shift the equilibrium position in the direction that produces fewer moles of gas. For example, increasing the pressure of a reaction that produces two moles of gas from one mole of gas will shift the equilibrium position to the left, favoring the formation of reactants.
• Catalyst: A catalyst is a substance that speeds up the rate of a chemical reaction without being consumed itself. A catalyst can shift the equilibrium position of a reaction, but it cannot change the extent to which the reaction reaches equilibrium.

Le Chatelier’s Principle

Le Chatelier’s principle states that when a stress is applied to a system at equilibrium, the system will respond in a way that tends to relieve the stress.

For example, if the concentration of a reactant is increased, the equilibrium position will shift to the right, favoring the formation of products. This is because the system is trying to relieve the stress of the increased concentration of reactants by consuming them and forming products.

Examples of Chemical Equilibrium

There are many examples of chemical equilibrium in the world around us. Here are a few:

• The evaporation of water is in equilibrium with its condensation.
• The dissolution of a salt in water is in equilibrium with its precipitation.
• The dissociation of a weak acid in water is in equilibrium with its association.
• The Haber process for the synthesis of ammonia is in equilibrium.

Importance of Chemical Equilibrium

Chemical equilibrium is an important concept in chemistry because it helps us to understand how chemical reactions work and how to control them. For example, chemical equilibrium is used in the production of many industrial chemicals and in the design of chemical reactors.

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