Chapter 6: Equilibrium Short Questions | chemistry Class 11 CBSE

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Chapter 6: Equilibrium Short Questions | chemistry Class 11 CBSE | ByteNotes.in

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Question 1

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What is dynamic equilibrium? How does it differ from static equilibrium?

In dynamic equilibrium, forward and reverse reactions occur simultaneously at equal rates, so no net change in concentrations is observed even though molecular-level processes continue.
In static equilibrium, all activity has stopped; in dynamic equilibrium, the system is actively exchanging species at the molecular level at the boundary between phases or between reactants and products.
Dynamic equilibrium is possible only in a closed system at a given temperature; measurable properties like concentration, vapour pressure, and colour remain constant at this stage.

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Question 2

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State the Law of Chemical Equilibrium (Law of Mass Action).

At a given temperature, the product of concentrations of reaction products raised to their stoichiometric coefficients divided by the product of concentrations of reactants raised to their stoichiometric coefficients has a constant value called the equilibrium constant Kc.
For the general reaction aA + bB ⇌ cC + dD, Kc = [C]c[D]d / [A]a[B]b, where the concentrations are equilibrium concentrations in mol L–1.
The equilibrium constant Kc is temperature-dependent and its value changes only when the temperature of the system changes; it is independent of initial concentrations.

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Question 3

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Write the relationship between Kp and Kc. Define all the terms used.

The relationship between Kp and Kc is given by Kp = Kc(RT)Δn, where R is the universal gas constant (0.0831 bar L mol–1 K–1) and T is the temperature in Kelvin.
Δn is the difference between the number of moles of gaseous products and the number of moles of gaseous reactants in the balanced chemical equation: Δn = (moles of gaseous products) – (moles of gaseous reactants).
When Δn = 0, Kp = Kc; when Δn > 0, Kp > Kc; and when Δn < 0, Kp < Kc. Pressure must be expressed in bar for this relation to hold since the standard state for pressure is 1 bar.

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Question 4

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What is the reaction quotient Qc? How is it used to predict the direction of a reaction?

The reaction quotient Qc is defined in the same way as the equilibrium constant Kc but uses concentrations measured at any time, not necessarily at equilibrium, for the reaction aA + bB ⇌ cC + dD: Qc = [C]c[D]d / [A]a[B]b.
If Qc < Kc, the concentrations of products are too low relative to reactants and the reaction proceeds in the forward direction to reach equilibrium.
If Qc > Kc, the concentrations of products are too high and the reaction proceeds in the reverse direction; if Qc = Kc, the system is already at equilibrium and no net reaction occurs.

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Question 5

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State Le Chatelier's principle. Give one example of its application in chemical industry.

Le Chatelier's principle states that if a system at equilibrium is subjected to a change in concentration, pressure, or temperature, the equilibrium shifts in such a direction so as to reduce or counteract the effect of the change.
In the Haber process for ammonia synthesis, N2(g) + 3H2(g) ⇌ 2NH3(g), since the forward reaction involves a decrease in moles of gas, increasing pressure shifts the equilibrium forward, increasing the yield of ammonia.
In practice, ammonia is continuously liquified and removed from the equilibrium mixture, maintaining Qc below Kc and driving the reaction forward; this is a direct industrial application of Le Chatelier's principle.

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Question 6

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Why are pure solids and liquids not included in the equilibrium constant expression?

The molar concentration of a pure solid or liquid is constant regardless of the amount present, because density and molar mass are fixed quantities for pure substances under given conditions.
Since [X(s)] and [X(l)] are constant, they are incorporated into the equilibrium constant itself, giving a modified constant that only includes the variable concentrations of gases and solutes.
For example, in CaCO3(s) ⇌ CaO(s) + CO2(g), the equilibrium constant is written as K'c = [CO2(g)] or Kp = pCO2, omitting the constant concentrations of the two solids.

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Question 7

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Explain the effect of a catalyst on a chemical equilibrium.

A catalyst increases the rate of both the forward and the reverse reactions by providing a new lower-energy pathway for the conversion of reactants to products and vice versa, lowering the activation energy by exactly the same amount for both directions.
A catalyst does not affect the equilibrium constant Kc or the equilibrium composition of the reaction mixture, since it increases both forward and reverse rates equally.
A catalyst allows the system to reach equilibrium faster but does not change the position of equilibrium; it also does not appear in the balanced chemical equation or the equilibrium constant expression.

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Question 8

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What is Henry's law? State its application in the context of gas-liquid equilibrium.

Henry's law 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.
This law governs the equilibrium CO2(gas) ⇌ CO2(in solution); soda water bottles are sealed under high CO2 pressure, and when opened, dissolved CO2 escapes to reach a new equilibrium corresponding to the lower partial pressure of CO2 in the atmosphere.
The solubility of a gas in a liquid decreases with increasing temperature, which is why heated soft drinks fizz more rapidly; practically, dissolved gas concentration in a liquid is proportional to the pressure (concentration) of the gas over the liquid.

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Question 9

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Define conjugate acid-base pair according to the Brønsted-Lowry theory.

According to Brønsted-Lowry theory, a conjugate acid-base pair consists of two species that differ only by one proton (H+); the species that donates the proton is the acid and the species that results after donation is the conjugate base.
For example, in NH3 + H2O ⇌ NH4+ + OH–, the conjugate acid-base pairs are: H2O (acid) and OH– (conjugate base), and NH3 (base) and NH4+ (conjugate acid).
A strong acid has a very weak conjugate base, and a weak acid has a relatively stronger conjugate base; the stronger the acid, the weaker the conjugate base, and vice versa.

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Question 10

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Distinguish between Arrhenius and Brønsted-Lowry definitions of acids and bases.

Arrhenius defines an acid as a substance that dissociates in water to give H+ ions and a base as a substance that produces OH– ions in water; this definition is limited to aqueous solutions.
Brønsted-Lowry defines an acid as a proton donor and a base as a proton acceptor; this is a broader definition applicable beyond aqueous solutions and explains the basicity of ammonia, which has no OH group.
A key difference is that Brønsted-Lowry theory introduces the concept of conjugate acid-base pairs and accounts for the amphoteric nature of water, which can act as both an acid (proton donor to NH3) and a base (proton acceptor from HCl).

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Question 11

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What is the ionic product of water (Kw)? Write its expression and state its value at 298 K.

Water undergoes autoionization: H2O(l) + H2O(l) ⇌ H3O+(aq) + OH–(aq); the equilibrium constant for this process, incorporating the constant concentration of water, is called the ionic product of water, Kw.
Kw = [H+][OH–]; experimentally, at 298 K, [H+] = [OH–] = 1.0 × 10–7 M, giving Kw = (1 × 10–7)2 = 1 × 10–14 M2.
Kw is temperature-dependent; it increases with temperature because the autoionization of water is endothermic. The relation pH + pOH = pKw = 14 holds at 298 K for all aqueous solutions.

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Question 12

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How does addition of an inert gas at constant volume affect equilibrium?

Addition of an inert gas at constant volume does not affect the equilibrium because the partial pressures of the reacting species and their molar concentrations remain unchanged.
The total pressure of the system increases, but since the inert gas does not participate in the reaction, neither the reaction quotient Qc nor the equilibrium constant Kc is affected.
The equilibrium remains undisturbed; in contrast, if the volume were allowed to increase (constant pressure) on addition of inert gas, the equilibrium would shift towards the side with more moles of gas.

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Question 13

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Define the term 'degree of ionization'. How does dilution affect the degree of ionization of a weak acid?

Degree of ionization (α) is defined as the fraction of the total initial amount of a weak electrolyte that has dissociated into ions at equilibrium; it can also be expressed as a percentage called percent dissociation.
For a weak acid HA, Ka = cα2/(1–α); when α is small, Ka ≈ cα2, giving α = √(Ka/c); since α is inversely proportional to √c, dilution (decreasing c) increases α.
Upon dilution, the degree of ionization of a weak acid increases, meaning a greater fraction of the acid molecules dissociate; however, the actual concentration of H+ ions decreases, causing the pH to increase.

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Chapter 6: Equilibrium | Short Questions

What is dynamic equilibrium? How does it differ from static equilibrium?

State the Law of Chemical Equilibrium (Law of Mass Action).

Write the relationship between Kp and Kc. Define all the terms used.

What is the reaction quotient Qc? How is it used to predict the direction of a reaction?

State Le Chatelier's principle. Give one example of its application in chemical industry.

Why are pure solids and liquids not included in the equilibrium constant expression?

Explain the effect of a catalyst on a chemical equilibrium.

What is Henry's law? State its application in the context of gas-liquid equilibrium.

Define conjugate acid-base pair according to the Brønsted-Lowry theory.

Distinguish between Arrhenius and Brønsted-Lowry definitions of acids and bases.

What is the ionic product of water (Kw)? Write its expression and state its value at 298 K.

How does addition of an inert gas at constant volume affect equilibrium?

Define the term 'degree of ionization'. How does dilution affect the degree of ionization of a weak acid?

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