Equilibrium MCQ Quiz - Objective Question with Answer for Equilibrium - Download Free PDF

Last updated on Jun 10, 2025

Latest Equilibrium MCQ Objective Questions

Equilibrium Question 1:

Match List I with List II.

  List - I
(Type of Salt)
  List - II
(Solubility Expression)
A. AB (e.g. AgCl) I. S = ∛(Ksp / 4)
B. AB2 (e.g. PbI2) II. S = √Ksp  
C. AB3 (e.g. Al(OH)3) III. S = (Ksp / 27)1/4
D. A3B2 (e.g. Ca3(PO4)2) IV. S = ∛(Ksp / 108)
    V. S = (Ksp / 27)1/3

Choose the correct answer from the options given below

  1. A - II, B - I, C - III, D - IV
  2. A - I, B - II, C - IV, D - III
  3. A - III, B - IV, C - II, D - I
  4. A - IV, B - III, C - I, D - II

Answer (Detailed Solution Below)

Option 1 : A - II, B - I, C - III, D - IV

Equilibrium Question 1 Detailed Solution

CONCEPT:

Solubility Product (Ksp) and Solubility (S)

  • The solubility product constant (Ksp) is the equilibrium constant for a solid substance dissolving in an aqueous solution.
  • For a salt dissolving into its constituent ions, the relation between solubility (S) and Ksp depends on the stoichiometry of the dissociation.
  • The general method is to:
    • Write the dissociation equation
    • Express ion concentrations in terms of solubility S
    • Substitute into the Ksp expression and solve for S

EXPLANATION:

  • A - II: AB type (e.g. AgCl)
    • Dissociation: AB ⇌ A+ + B
    • Ksp = S × S = S² ⇒ S = √Ksp
  • B - I: AB2 type (e.g. PbI2)
    • Dissociation: AB2 ⇌ A2+ + 2B
    • Ksp = (2S)² × S = 4S³ ⇒ S = ∛(Ksp/4)
  • C - III: AB3 type (e.g. Al(OH)3)
    • Dissociation: AB3 ⇌ A3+ + 3B
    • Ksp = S × (3S)³ = 27S⁴ ⇒ S = (Ksp/27)1/4
  • D - IV: A3B2 type (e.g. Ca3(PO4)2)
    • Dissociation: A3B2 ⇌ 3A2+ + 2B3−
    • Ksp = (3S)³ × (2S)² = 108S⁵ ⇒ S = ∛(Ksp/108)

Therefore, the correct answer is: Option 1) A - I, B - II, C - III, D - IV

Equilibrium Question 2:

Higher yield of NO in
N₂(g) + O₂ → 2NO(g) can be obtained at
[ΔH of the reaction = +180.7 kJ mol⁻¹]

A. higher temperature

B. lower temperature

C. higher concentration of N₂

D. higher concentration of O₂

Choose the correct answer from the options given below:

  1. A, D only
  2. B, C only
  3. B, C, D only
  4. A, C, D only

Answer (Detailed Solution Below)

Option 4 : A, C, D only

Equilibrium Question 2 Detailed Solution

CONCEPT:

Effect of Temperature and Concentration on Chemical Equilibrium

  • The given reaction is:

    N₂(g) + O₂(g) → 2NO(g)

  • This reaction is endothermic because ΔH = +180.7 kJ mol⁻¹.
  • According to Le Chatelier's Principle:
    • For an endothermic reaction (ΔH > 0), increasing the temperature shifts the equilibrium toward the products (higher yield of NO).
    • Increasing the concentration of reactants (N₂ and O₂) also shifts the equilibrium toward the products, increasing the yield of NO.

EXPLANATION:

  • Higher temperature: Since the reaction is endothermic, increasing the temperature supplies more energy to drive the reaction forward.
  • Higher concentration of N₂: Adding more N₂ shifts the equilibrium toward the products.
  • Higher concentration of O₂: Adding more O₂ also shifts the equilibrium toward the products.

Therefore, the correct answer is A, C, D only.

Equilibrium Question 3:

For the reaction A(g) ⇌ 2B(g), the backward reaction rate constant is higher than the forward reaction rate constant by a factor of 2500, at 1000 K.
[Given: R = 0.0831 L atm mol⁻¹ K⁻¹]
Kₚ for the reaction at 1000 K is

  1. 83.1
  2. 2.077 × 10⁵
  3. 0.033
  4. 0.021

Answer (Detailed Solution Below)

Option 3 : 0.033

Equilibrium Question 3 Detailed Solution

CONCEPT:

Relation Between KC and KP

  • The equilibrium constant Kp is related to Kc by the equation:

    Kp = KC (RT)Δng

  • Where:
    • KC is the equilibrium constant in terms of concentration.
    • R is the gas constant (0.0831 L atm mol-1 K-1).
    • T is the temperature in Kelvin (1000 K in this case).
    • Δng is the change in the number of moles of gas (products - reactants).
  • For the given reaction, the backward rate constant is 2500 times the forward rate constant, and we are asked to calculate Kp at 1000 K.

EXPLANATION:

  • The relationship between the rate constants for the forward and backward reactions is:

    \(KC = \frac{k{\text{f}}}{k{\text{b}}}\)

    Given that the backward rate constant is 2500 times the forward rate constant, we have:

    \(KC = \frac{k{\text{f}}}{2500 k{\text{f}}} = \frac{1}{2500}\)

  • Next, we use the equation to find Kp :

    Kp = KC (RT)Δng

    Given that Δng = 1 (because there is a change in the number of moles from 1 mole of reactants to 2 moles of products), we substitute the known values:

    \(Kp = \frac{1}{2500} (0.0831 × 1000)^1\)

  • Solving the equation:

    \(Kp = \frac{1}{2500} \times 83.1 = 0.033\)

Therefore, the correct value of Kp is 0.033.

Equilibrium Question 4:

Phosphoric acid ionizes in three steps with their ionization constant values Kₐ₁, Kₐ and Kₐ respectively, while K is the overall ionization constant. Which of the following statements are true?
A. log K = log Kₐ₁ + log Kₐ₂ + log Kₐ₃
B. H₃PO₄ is stronger acid than H₂PO₄⁻ and HPO₄²⁻.
C. Kₐ₁ > Kₐ₂ > Kₐ₃
D.
Choose the correct answer from the options below:

  1. A and B only
  2. A and C only
  3. B, C and D only
  4. A, B and C only

Answer (Detailed Solution Below)

Option 4 : A, B and C only

Equilibrium Question 4 Detailed Solution

CONCEPT:

Phosphoric Acid Ionization and Acid Strength

  • Phosphoric acid (H3PO4) ionizes in three steps, leading to the formation of H2PO4-, HPO42-, and PO43- respectively.
  • Each ionization step has an associated ionization constant, denoted as Kₐ₁, Kₐ₂, and Kₐ₃.
  • The overall ionization constant (K) is the product of the individual ionization constants: K = Kₐ₁ × Kₐ₂ × Kₐ₃.
  • The logarithm of the overall ionization constant is the sum of the logarithms of the individual constants: log K = log Kₐ₁ + log Kₐ₂ + log Kₐ₃.
  • The strength of an acid decreases as it ionizes further (i.e., H₃PO₄ is stronger than H₂PO₄-, which is stronger than HPO₄2-).

EXPLANATION:

  • Statement A: log K = log Kₐ₁ + log Kₐ₂ + log Kₐ₃ is correct because the overall ionization constant is the product of the individual ionization constants, and the logarithm of a product is the sum of the logarithms of the individual factors.
  • Statement B: H₃PO₄ is stronger than H₂PO₄- and HPO₄2- is correct because the acid strength decreases with successive ionizations. The first ionization constant (Kₐ₁) is the largest, making H₃PO₄ the strongest acid among its ionized forms.
  • Statement C: Kₐ₁ > Kₐ₂ > Kₐ₃ is correct because the ionization constants decrease sequentially as the molecule loses protons, reflecting decreasing acid strength with each step.
  • Statement D: K = (Kₐ₁ + Kₐ₂)/2 is incorrect because the overall ionization constant is the product of Kₐ₁, Kₐ₂, and Kₐ₃, not an average of the first two constants.

Therefore, the correct answer is A, B, and C only.

Equilibrium Question 5:

Comprehension:

Molarity is number of moles of solute dissolved per litre of the solution while normality is number of gm-equivalent of solute dissolved per litre of solution. Molality is number of moles of solute dissolved per Kg of solvent. Normality and molarity change with change of temperature of solution but molality is independent of temperature. In case of monobasic acid normality and molarity are equal but in dibasic acid or base molarity is two times of normality. In redox and neutralisation processes number of milliequivalents of reactants as well as products are always equal

The volume of 0.1 M Ca(OH)2 required to neutralise 0.2 M H3PO3 solution of volume 0.25 dm3 will be (in mL)

Answer (Detailed Solution Below) 500

Equilibrium Question 5 Detailed Solution

CONCEPT:

Neutralisation Reaction using N1V1 = N2V2 Method

  • In acid-base neutralisation, the milliequivalents of acid = milliequivalents of base.
  • The formula used is:

    Nacid × Vacid = Nbase × Vbase

  • Normality (N) = Molarity × n-factor
  • For H3PO3 (phosphorous acid), only 2 hydrogen ions are ionisable (n-factor = 2).
  • For Ca(OH)2, each molecule provides 2 OH ions (n-factor = 2).

EXPLANATION:

  • Given:
    • Molarity of H3PO3 = 0.2 M
    • Volume of H3PO3 = 0.25 L
    • Molarity of Ca(OH)2 = 0.1 M
  • Step 1: Calculate Normalities
    • Nacid = 0.2 × 2 = 0.4 N
    • Nbase = 0.1 × 2 = 0.2 N
  • Step 2: Apply formula:

    Nacid × Vacid = Nbase × Vbase

    0.4 × 250 mL = 0.2 × Vbase

    Vbase = (0.4 × 250) / 0.2 = 500 mL

Therefore, the volume of Ca(OH)2 required to neutralise the acid is 500 mL.

Top Equilibrium MCQ Objective Questions

Which of the following statements are correct about an aqueous solution of an acid and a base?

1. Higher the pH, stronger is the acid.

2. Higher the pH, weaker is the acid.

3. Lower the pH, stronger is the base.

4. Lower the pH, weaker is the base.

  1. 1 and 3
  2. 2 and 3
  3. 1 and 4
  4. 2 and 4

Answer (Detailed Solution Below)

Option 4 : 2 and 4

Equilibrium Question 6 Detailed Solution

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Concept:

The pH Scale

  • The pH is the measure of the concentration of acid in the solution. 
  • It is measured from the scale of 1 to 14. 
  • If the solution is acidic or basic, depends upon the following parameters. 
  • pH < 7 - Acidic
  • pH > 7 - Basic
  • pH = 7 - Neutral

​Conclusion:

  • So, from the above explanation, it is clear that the Higher the pH, the weaker is the acid, and the Lower the pH, the weaker is the base.
  • So, Statements 2 and 4 are correct. 

Additional Information

pH VALUE

EXAMPLES

0

battery acid

1

gastric acid

2

lemon juice, vinegar

3

orange juice, soda

4

tomato juice, acid rain

5

black coffee, bananas

6

urine, milk

7

pure water, neutral solution

8

seawater, eggs

9

baking soda

10

milk of magnesia

11

ammonia solution

12

soapy water

13

bleach, oven cleaner

14

liquid drain cleaner

The pH of water at 25°c is 7. When it is heated to 100°C, the pH of water ________

  1. Increases
  2. Decreases
  3. Remains same
  4. Decreases up to 27°C and then increases 

Answer (Detailed Solution Below)

Option 2 : Decreases

Equilibrium Question 7 Detailed Solution

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  • The pH of water at 25°c is 7. When it is heated to 100°C, the pH of water decreases to 6.14.
  • On increasing the temperature of the water, the equilibrium would again move to lower the temperature.
  • It would happen by absorbing the extra heat. 
  • Hydrogen ions and hydroxide ions will remain at the same concentration in pure water and the water remains neutral even if its pH changes.

The pH of the solution containing 50 mL each of 0.10 M sodium acetate and 0.01 M acetic acid is [Given pKa of CHCOOH = 4.57]

  1. 2.57
  2. 5.57
  3. 3.57
  4. 4.57

Answer (Detailed Solution Below)

Option 2 : 5.57

Equilibrium Question 8 Detailed Solution

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Concept:

pH of acidic buffer - 

  • Buffer solution - Buffer solution is a mixture of a weak acid and its conjugate base or a weak base and its conjugate acid. It is a solution highly known for maintaining its pH on dilution with a small amount of strong acid or strong base by maintaining the H+ ion concentration inside it. 
  • Two types of buffer solutions are there
  • Acidic buffer - Highly specific for pH<7 and is a mixture of a weak acid and its conjugate base.
  • Weak buffer - Highly specific for pH>7 and is a mixture of a weak base and its conjugate acid.
  • pH of the acidic buffer is given by the formula -

pH = pKa + log ([salt]/[acid])

where pK= negative logarithms of the acid dissociation constant.

As acetic acid is a weak acid and sodium acetate is its conjugate salt with a strong base(NaOH), they form an acidic buffer whose pH is calculated by the formula -

pH = pKa + log ([salt]/[acid])

Chemical reaction - CH3COOH + NaOH \(\rightarrow\) CH3COONa +H2O

Calculation -

Given,

  • pKa = 4.57 (negative logarithms of the acid dissociation constant of acetic acid)
  • [acid] = 0.01 M (concentration of acetic acid)
  • [salt] = 0.10 M (concentration of salt)

Put all these values in the formula,

pH = pKa + log ([salt]/[acid])

=4.57 + log(0.10 / 0.01)

=4.57 + log(10)

= 4.57 + 1   -- (∵ log10 = 1)

=5.57

So, the pH of the buffer solution is = 5.57

Hence, the correct answer is option 2.

Which of the following are equal for a chemical system at equilibrium?

  1. The concentration of reactant and products
  2. The rate of the forward and reverse reactions
  3. The rate constants for the forward and reverse reactions
  4. Both 2 & 3

Answer (Detailed Solution Below)

Option 2 : The rate of the forward and reverse reactions

Equilibrium Question 9 Detailed Solution

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Chemical Equilibrium: Consider a general case of a reversible reaction

A + B ⇔ C + D

F1 Madhuri Engineering 25.04.2022 D1

  • With respect of time, the rate of forward reaction decreased, and the rate of backward reaction increased.
  • With an instant of time, a stage is reached at which the rate of forward and reverse reaction become equal and the concentration of reactant and product becomes constant.
  • The equilibrium is dynamic in 

Who synthesised acetic acid for the first time?

  1. Berthelot
  2. F Wohler
  3. Berzelius
  4. Kolbe

Answer (Detailed Solution Below)

Option 4 : Kolbe

Equilibrium Question 10 Detailed Solution

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The correct answer is Kolbe.

Key Points

  • Synthesis of Acetic Acid
    • Acetic acid was first made or created from inorganic components in 1845 by German scientist Hermann Kolbe.
    • He created carbon disulfide with chlorine, and the end result was carbon tetrachloride.
    • After that, tetrachloroethylene was produced through thermal decomposition, trichloroacetic acid was created through aqueous chlorination, and acetic acid was finally produced through an electrolytic reduction.
    • He exclusively used inorganic materials made of hydrocarbons.
    • Carbon, hydrogen, and oxygen are inorganic elements that are used to create acetic acid while acetic acid is an organic one.
    • Acetic acid has a chemical formula and is a weak acid.
    • Kolbe's work was groundbreaking since he synthesised organic chemicals from inorganic ones.
    • In addition to bacterial fermentation, acetic acid is also made synthetically.
    • Some fruits naturally contain it.

Additional Information

  • Kolbe's work was groundbreaking since he synthesised organic chemicals from inorganic ones
  • In addition to bacterial fermentation, acetic acid is also made synthetically.
  •  some fruits naturally contain it.
Scientist Creation Year
Berthelot Methane 1856
F Wohler Urea 1828
Berzelius

Cerium

selenium 

1803

1817

What is the pH value of a salt made up of a strong acid and a weak base?

  1. More than 7
  2. Less than 7
  3. Between 10 to 14
  4. More than one of the above
  5. None of the above

Answer (Detailed Solution Below)

Option 2 : Less than 7

Equilibrium Question 11 Detailed Solution

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The correct answer is Less than 7

Key Points 

  • When a salt is formed from the neutralization of a strong acid and a weak base, the resulting solution is acidic. This is because the strong acid completely dissociates in water, providing a high concentration of (H+) ions, while the weak base does not dissociate completely and therefore provides fewer (OH-) ions to neutralize the (H+). As a result, the (H+) concentration in the solution is higher than the (OH-) concentration, leading to an acidic solution.

Therefore, the pH value of a salt made up of a strong acid and a weak base is Less than 7

  • The cations of strong acids do not hydrolyze and therefore the solutions of salts formed by strong acids and strong bases are neutral, i.e. their pH is 7, whereas the solutions of salts formed by strong bases and weak acids are alkaline, i.e. their pH>7.
  • Salts that are from strong bases and weak acids do hydrolyze, which gives them a pH greater than 7.
  • The pH value range of a weak base is 7.1 to 10 whereas a strong base is 10.1 to 14.
  • All substances can be classified as neutral (with a pH of about 7), basic (pH greater than 7), or acidic (pH less than 7), and the pH tells us how strong or weak that substance is as well. For example, a substance with pH = 8 is a very weak base, but a substance with pH = 3 is a strong acid.

Which pair of natural source - acid is correct?

I. Tamarind - oxalic acid

II. Yogurt - lactic acid

  1. Nor neither II
  2. II Only 
  3. Both I and II
  4. । Only

Answer (Detailed Solution Below)

Option 2 : II Only 

Equilibrium Question 12 Detailed Solution

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The correct answer is II Only.

Key Points List of Acid and Natural Resources:

SN  Acid  Natural Resource
1 Acetic acid  Vinegar
2 Malic Acid   Apple
3 Tartaric Acid Grapes
4 citric acid Citrus fruit 
5 Lactic acid Milk and Yogurt 
6 Benzoic acid berries
7 Formic Acid Scorpion and Ant 
8 Citric acid Lemon
9 Nitric acid Alum 
10 Sulfuric acid Green Vitriol

 

Consider the following reversible chemical reactions:

\({A_2}\left( g \right)\; + \;{B_2}\left( g \right)\begin{array}{*{20}{c}} {{K_1}}\\ \rightleftharpoons \end{array}\;2AB\left( g \right)\)      …. (1)

\(6AB\left( g \right)\;\begin{array}{*{20}{c}} {{K_2}}\\ \rightleftharpoons \end{array}3{A_2}\left( g \right) + 3{B_2}\left( g \right)\)        …. (2)

The relation between K1 and K2 is:

  1. K1 × K2 = 1/3
  2. K2 = K1+3
  3.  K2 = K1-3
  4. K1 × K2 = 3

Answer (Detailed Solution Below)

Option 3 :  K2 = K1-3

Equilibrium Question 13 Detailed Solution

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Explanation:

A reversible reaction is a chemical reaction where the reactants form products that, in turn, react together to give the reactants back. Reversible reactions will reach an equilibrium point where the concentrations of the reactants and products will no longer change.

According to equilibrium constant, Kc

\({{\rm{K}}_{\rm{c}}} = \frac{{{{\left[ {{\rm{product}}} \right]}^{\rm{m}}}}}{{{{\left[ {{\rm{reactant}}} \right]}^{\rm{n}}}}}\)

Consider the given first equation:

\({A_2}\left( g \right)\; + \;{B_2}\;\left( g \right)\begin{array}{*{20}{c}} {{K_1}}\\ \rightleftharpoons \end{array}\;2AB\left( g \right)\)

\({K_1} = \frac{{{{\left[ {AB} \right]}^2}}}{{\left[ {{A_2}} \right]\left[ {{B_2}} \right]}}\)

Consider the given second equation:

\(6AB\left( g \right)\begin{array}{*{20}{c}} {{K_2}}\\ \rightleftharpoons \end{array}\;3{A_2}\left( g \right) + 3{B_2}\left( g \right)\)

\({K_2} = \frac{{{{\left[ {{A_2}} \right]}^3}{{\left[ {{B_2}} \right]}^3}}}{{{{\left[ {AB} \right]}^6}}}\)

\(= \frac{1}{{{{\left( {\frac{{{{\left[ {AB} \right]}^2}}}{{\left[ {{A_2}} \right]\left[ {{B_2}} \right]}}} \right)}^3}}}\)

\(= \frac{1}{{K_1^3}}\)

 → K2 = K1-3

Which of the following pairs is Lewis acids?

  1. NH4+, BCl3
  2. BCl3, FeCl3
  3. AlCl3, H2CO3
  4. HCl, HNO3

Answer (Detailed Solution Below)

Option 2 : BCl3, FeCl3

Equilibrium Question 14 Detailed Solution

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 Concept:

Lewis concept of acids and bases is given as follows:

Lewis acid Lewis base
These are the species hey which have electron deficiency hey in them. These are the species that have an electron-rich density in them.
They don't have their octet fulfilled. They have excessive electrons.
They can accept an electron pair. They can easily donate a lone pair.
Are usually electrophilic in nature. They're basically nucleophilic in nature.

Lewis acids can be positive or neutral.

Examples are H+, H3O+ , CH3+ , NO+ , AlCl3 etc.

Lewis acids can be negative or neutral.

Examples are Cl-, F-, CN-, H-O-H, R-O-H, NO2-.

 

The interaction between Lewis acid and base- 

  • Lewis acid has a vacant unoccupied orbital called LUMO.
  • Lewis bases have fulfilled occupied orbital called HOMO.
  • Electron density from HOMO is donated to LUMO of Lewis acid.
  • A simple HOMO and LUMO interaction is shown below:

F1 Shraddha Pooja J 11.02.2021 D1

This donation of electron density forms a covalent coordinate between the acid and the base,

Explanation:

BCl3 -

  • BCl3 molecule has an empty p orbital because its octet is not fulfilled.
  • It can accept an electron pair inti this orbital

Thus it electron-deficient and a Lewis acid

F1 Shraddha Pooja J 11.02.2021 D2

FeCl3:

  • Fe is a transition metal having empty d orbitals.
  • It can easily accept a lone pair in the empty 4d orbital and can act as Lewis acid.

F1 Shraddha Pooja J 11.02.2021 D3

Hence, BCl3, FeCl3 is a pair of Lewis acids.

Additional Information

  • AlCl3: AlClalso has empty p orbitals and thus is an electrophile.

 

F1 Shraddha Pooja J 11.02.2021 D4

  • Proton acceptors are Bronsted-Lowry bases, while Proton donors are Bronsted-Lowry acids

Thus, HCl, HNO3 and H2CO3 NH4+ are Bronsted acids.

N2 + 3H2 ⇋ 2NH3 + Heat

The Le Chatelier's principle suggests that are required to drive the reaction to the right and thus form NH3

  1. high pressure and low temperature
  2. low pressure and low temperature
  3. high pressure and high temperature
  4. low pressure and high temperature

Answer (Detailed Solution Below)

Option 1 : high pressure and low temperature

Equilibrium Question 15 Detailed Solution

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The correct answer is high pressure and low temperature.

Key Points

  • This equation is an example of Equilibrium.
  • Equilibrium is a condition that occurs when a chemical reaction is reversible, and the forward and reverse reactions occur simultaneously, at the same rate.
  • Le Chatlier's Principle- "if a chemical system at equilibrium experiences a change in concentration, temperature or total pressure, the equilibrium will shift in order to minimize that change" 
  • Any decrease in nitrogen or hydrogen pulls the reaction towards the left side.
  • Any decrease in ammonia or temperature pulls the reaction towards the right side.
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