Short Q&A for Chemical and Ionic Equilibrium (Class 11 NCERT)

Over 110 Short Q&A for Chemical and Ionic Equilibrium (Class 11 NCERT)

Part 1: Chemical Equilibrium (General Concepts)

Q. No.	Question	Answer
1	Define a Reversible Reaction.	A reaction that can proceed in both forward and reverse directions.
2	What is a State of Chemical Equilibrium?	The state where the rate of the forward reaction equals the rate of the reverse reaction.
3	Is chemical equilibrium static or dynamic?	Dynamic (reactions continue, but concentrations remain constant).
4	How is the Equilibrium Constant (K) expressed in terms of concentration?	Kc​ (Concentration in mol L−1).
5	How is the Equilibrium Constant (K) expressed in terms of partial pressure?	Kp​ (Pressure in bar or atm).
6	Write the relationship between Kp​ and Kc​.	Kp​=Kc​(RT)Δng​.
7	What does Δng​ represent in the Kp​ and Kc​ relationship?	(Moles of gaseous products) - (Moles of gaseous reactants).
8	When is Kp​=Kc​?	When Δng​=0.
9	Does the magnitude of K indicate the speed of a reaction?	No, K indicates the extent of the reaction, not the speed.
10	What does a large value of K (≈103) indicate?	The reaction proceeds almost to completion (products are favoured).
11	What does a very small value of K (≈10−3) indicate?	The reaction hardly proceeds (reactants are favoured).
12	What is the Reaction Quotient (Q)?	The expression for Kc​ evaluated at any point during the reaction, not necessarily at equilibrium.
13	If Qc​<Kc​, which direction will the reaction proceed?	Forward direction (to increase products and reach equilibrium).
14	If Qc​>Kc​, which direction will the reaction proceed?	Reverse direction (to decrease products and reach equilibrium).
15	State Le Chatelier's Principle.	When a system at equilibrium is subjected to a change in concentration, temperature, or pressure, the system shifts to counteract the change.
16	How does increasing the concentration of a reactant affect equilibrium?	The equilibrium shifts in the forward direction.
17	How does increasing the pressure affect equilibrium (for Δng​=0)?	It shifts towards the side with the fewer number of gaseous moles.
18	For an exothermic reaction, how does increasing the temperature affect K?	K decreases (equilibrium shifts in the reverse direction).
19	For an endothermic reaction, how does increasing the temperature affect K?	K increases (equilibrium shifts in the forward direction).
20	Does adding an inert gas at constant volume affect equilibrium?	No, it only changes total pressure, not partial pressures.
21	Does adding a catalyst affect the value of K?	No, a catalyst only helps attain equilibrium faster.
22	What is a Homogeneous Equilibrium?	All reactants and products are in the same phase (e.g., all gases or all liquids).
23	What is a Heterogeneous Equilibrium?	Reactants and products are in different phases (e.g., solid and gas).
24	How are pure solids and liquids treated in the K expression?	Their concentration is taken as unity (1).

Part 2: Ionic Equilibrium (Acids, Bases, and Salts)

Q. No.	Question	Answer
25	What is an Electrolyte?	A substance that dissociates into ions in an aqueous solution.
26	Name the three theories for acids and bases.	Arrhenius, Brønsted-Lowry, and Lewis.
27	Define an Arrhenius Acid.	A substance that produces H+ ions in water.
28	Define an Arrhenius Base.	A substance that produces OH− ions in water.
29	Define a Brønsted-Lowry Acid.	A substance that is a proton (H+) donor.
30	Define a Brønsted-Lowry Base.	A substance that is a proton (H+) acceptor.
31	What is a Conjugate Acid-Base Pair?	A pair that differs by only one proton (H+).
32	What is the conjugate acid of NH3​?	NH4+​ (Ammonium ion).
33	What is the conjugate base of H2​O?	OH− (Hydroxide ion).
34	What is an Amphoteric Substance?	A substance that can act as both an acid and a base (e.g., H2​O).
35	Define a Lewis Acid.	An electron pair acceptor.
36	Give one example of a Lewis Acid.	BF3​ or AlCl3​ (electron deficient molecules).
37	Define a Lewis Base.	An electron pair donor.
38	Give one example of a Lewis Base.	NH3​ or H2​O (molecules with lone pairs).
39	What is the significance of the ionization constant (Ka​ or Kb​)?	It measures the strength of an acid (Ka​) or base (Kb​).
40	What is the relationship between pKa​ and acid strength?	Smaller the pKa​, stronger the acid.
41	What is the Ionization Constant of Water (Kw​)?	Kw​=[H3​O+][OH−].
42	What is the value of Kw​ at 298 K?	1.0×10−14 (mol2 L−2).
43	What is the relationship between pKa​,pKb​, and pKw​?	pKa​+pKb​=pKw​ (=14 at 298 K).
44	Define pH scale.	pH=−log10​[H+] (or [H3​O+]).
45	What is the pH of a neutral solution at 298 K?	7
46	What is the pOH of a solution with [OH−]=10−4 M?	4 (pOH=−log(10−4)).
47	What is the pH of the solution in the previous question?	10 (pH+pOH=14).
48	What is the minimum condition for an acid to be considered a strong acid?	It must be almost completely ionized in an aqueous solution.
49	What is the Ostwald's Dilution Law used for?	To calculate the degree of dissociation (α) of a weak electrolyte.
50	What is Hydrolysis of a Salt?	The reaction of the salt's ion (cation or anion) with water to produce acid and base.
51	What kind of salt produces an acidic solution upon hydrolysis?	Salt of a strong acid and weak base (e.g., NH4​Cl).
52	What kind of salt produces a basic solution upon hydrolysis?	Salt of a weak acid and strong base (e.g., CH3​COONa).
53	What kind of salt produces a neutral solution upon hydrolysis?	Salt of a strong acid and strong base (e.g., NaCl).

Part 3: Buffer Solutions and Solubility Product

Q. No.	Question	Answer
54	Define a Buffer Solution.	A solution that resists change in its pH upon the addition of small amounts of acid or base.
55	What is a Simple Buffer?	A solution of a salt of a weak acid and weak base (rarely used).
56	What is an Acidic Buffer composed of?	A weak acid and its salt with a strong base (e.g., CH3​COOH+CH3​COONa).
57	What is a Basic Buffer composed of?	A weak base and its salt with a strong acid (e.g., NH3​+NH4​Cl).
58	Write the Henderson-Hasselbalch equation for an acidic buffer.	pH=pKa​+log([Salt]/[Acid]).
59	Write the Henderson-Hasselbalch equation for a basic buffer.	pOH=pKb​+log([Salt]/[Base]).
60	What is the pH of an acidic buffer when [Salt]=[Acid]?	pH=pKa​.
61	Define Solubility Product Constant (Ksp​).	The equilibrium constant for the equilibrium between a sparingly soluble ionic solid and its ions in a saturated solution.
62	If the ionic product (Qsp​) is less than Ksp​, what is the state of the solution?	Unsaturated (more salt can dissolve).
63	If the ionic product (Qsp​) is greater than Ksp​, what happens?	Precipitation occurs (supersaturated).
64	Write the Ksp​ expression for AgCl.	Ksp​=[Ag+][Cl−].
65	Write the Ksp​ expression for BaSO4​.	Ksp​=[Ba2+][SO42−​].
66	Write the Ksp​ expression for CaF2​.	Ksp​=[Ca2+][F−]2.
67	What is the effect of the Common Ion Effect on the solubility of a sparingly soluble salt?	The solubility decreases.
68	Why is HCl gas passed into a solution before precipitating Group II metal sulphides?	To suppress the ionization of H2​S (common ion H+) and control [S2−] to precipitate only Group II sulphides.
69	What is the relationship between solubility (s) and Ksp​ for AgCl?	Ksp​=s2.
70	What is the relationship between solubility (s) and Ksp​ for CaF2​?	Ksp​=4s3 (s×(2s)2).

Part 4: Numerical and Conceptual Questions (Mixed)

Q. No.	Question	Answer
71	If K for a reaction is 4×1020, what is the sign of ΔG∘?	Negative (ΔG∘=−RTlnK).
72	What is the effect of a catalyst on the activation energy of the forward reaction (Eaf​)?	Eaf​ decreases.
73	What is the pH of 10−8 M HCl?	6.96 (Must consider H+ from water due to low concentration).
74	In the reaction H2​+I2​⇌2HI, $\Delta n_g = $?	0 (2−(1+1)).
75	In the reaction PCl5​(g)⇌PCl3​(g)+Cl2​(g), how does increasing pressure affect PCl5​ amount?	PCl5​ amount increases (shifts to fewer moles).
76	In the reaction N2​(g)+3H2​(g)⇌2NH3​(g)(ΔH<0), how does increasing T affect NH3​ yield?	NH3​ yield decreases (shifts to reverse for exothermic reaction).
77	In a buffer solution, which component consumes the added acid (H+)?	The conjugate base (salt component).
78	Why is the pH of NaCl solution 7?	Neither Na+ nor Cl− undergo hydrolysis.
79	Which one is the weaker acid: one with pKa​=5 or one with pKa​=3?	The acid with pKa​=5 (higher pKa​ means weaker acid).
80	Is O2− a strong or weak Brønsted base?	A strong Brønsted base (readily accepts H+ to form OH−).
81	What is the relationship between the strength of an acid and its conjugate base?	A strong acid has a weak conjugate base, and vice versa.
82	How does dilution affect the degree of ionization (α) of a weak acid?	α increases (Ostwald's Law).
83	What does Ksp​ really represent in terms of concentration?	The maximum product of ion concentrations in a saturated solution.
84	What is the main reason for the Common Ion Effect?	The shift in equilibrium due to the addition of an ion already present in the system (Le Chatelier's Principle).
85	Why must pure water ionization always be considered for very dilute strong acid/base solutions?	The concentration of H+ from water becomes comparable to the added acid/base concentration.

Part 5: More Definitions and Concepts

Q. No.	Question	Answer
86	If an overall reaction is the sum of two steps, how is the overall K related to K1​ and K2​?	Koverall​=K1​×K2​.
87	If an equilibrium reaction is reversed, how does the new K′ relate to the original K?	K′=1/K.
88	If an equilibrium reaction is doubled, how does the new K′′ relate to the original K?	K′′=K2.
89	What is the value of pKw​ at 298 K?	14 (pKw​=−logKw​).
90	What is the common ion in the CH3​COOH/CH3​COONa buffer?	CH3​COO− (Acetate ion).
91	What is the effect of pressure on solubility of gases in liquids (Henry's Law)?	Solubility increases with increasing pressure.
92	Define Buffer Capacity.	The amount of acid or base a buffer can absorb before its pH changes significantly.
93	What is the pH of a solution when [H+]=[OH−]?	Neutral pH (usually 7).
94	What is the relationship between the degree of hydrolysis (h) and Kh​ for a weak acid/strong base salt?	Kh​=Ch2 (C is concentration).
95	Write the expression for the hydrolysis constant (Kh​) of NH4​Cl.	Kh​=Kw​/Kb​.
96	Write the expression for the hydrolysis constant (Kh​) of CH3​COONa.	Kh​=Kw​/Ka​.
97	What is the primary species in water responsible for making a solution basic?	OH− (Hydroxide ion).
98	What is the primary species in water responsible for making a solution acidic?	H3​O+ (Hydronium ion, or simplified H+).
99	What is the minimum pH attainable by an aqueous solution?	Theoretically 0 or less (for highly concentrated strong acids).
100	Why must Ksp​ always be determined in a saturated solution?	Because Ksp​ is an equilibrium constant.
101	Does adding an inert gas at constant pressure affect equilibrium?	Yes, it increases volume, decreasing concentration/pressure of reactants, shifting equilibrium to the side with more gaseous moles.
102	What is the maximum pH attainable by an aqueous solution?	Theoretically 14 or more (for highly concentrated strong bases).
103	What is the relationship between ΔG∘ and K at equilibrium?	ΔG∘=−RTlnK (Standard Gibbs Free Energy equals zero at equilibrium).
104	Which theory is the most general for acids and bases?	Lewis Theory (includes species without H+ or OH−).
105	Does the value of Kw​ depend on temperature?	Yes, Kw​ increases with temperature (ionization of water is endothermic).
106	How does the solubility of an ionic solid typically change with increasing temperature?	Solubility increases (most dissolution processes are endothermic).
107	What is the common ion used to precipitate Group III hydroxides (Al3+,Fe3+)?	OH− (from NH4​OH and NH4​Cl buffer).
108	Is BF3​ a better Lewis acid or a Brønsted acid?	A Lewis Acid (electron deficient, lacks H+).
109	What type of salt solution requires both Ka​ and Kb​ of its parent acid/base to determine pH?	Salt of a weak acid and weak base.
110	If pKa​=pKb​ for the parent acid/base of a salt, what is the pH of the salt solution?	7 (Neutral).
111	What is the minimum Ka​ an acid must have to be practically considered a weak acid?	Ka​<1 (typically 10−3 to 10−10).
112	What is the value of ΔG when a reaction is at equilibrium?	ΔG=0.