The following equilibrium is established when hydrogen chloride is dissolved in acetic acid:

$$\(\small {\mathrm{HCl}+\mathrm{CH}_3 \mathrm{COOH} \rightleftharpoons \mathrm{Cl}^{-}+\mathrm{CH}_3 \mathrm{COOH}_2^{+}}\)

The set that characterizes the conjugate acid-base pair is:

1.	\(\small {(HCl, CH_3COOH) ~\text {and}~ (CH_3COOH_2^+ , Cl^-)}\)
2.	\(\small { (HCl, CH_3COOH_2^+)~ \text {and} ~(CH_3COOH, Cl^-)}\)
3.	\(\small {(CH_3COOH^+_2 , HCl)~ \text {and} ~(Cl^-, CH_3COOH) }\)
4.	\(\small {(HCl, Cl^-) \text {and}~(CH_3COOH_2^+ , CH_3COOH)}\)

Which of the following is an example of a reversible reaction?

A certain weak acid has a dissociation constant of $1.0\times {10}^{-4}.$ The equilibrium constant for its reaction with a strong base is

1. $1.0\times {10}^{10}$

2. $1.0\times {10}^{14}$

3. $1.0\times {10}^{-10}$

4. $1.0\times {10}^{-4}$

Find the condition under which CaF₂ (Ksp = 1.7 × 10⁻¹⁰) precipitate is formed
when equal volumes of two solutions are mixed.

1. ${10}^{-2}<mspace\; width="1em"></mspace>\mathrm{M}<mspace\; width="1em"></mspace>{\mathrm{Ca}}^{2+}+{10}^{-3}<mspace\; width="1em"></mspace>{\mathrm{MF}}^{-}$

2. ${10}^{-3}<mspace\; width="1em"></mspace>\mathrm{M}<mspace\; width="1em"></mspace>{\mathrm{Ca}}^{2+}+{10}^{-5}<mspace\; width="1em"></mspace>{\mathrm{MF}}^{-}$

3. ${10}^{-4}<mspace\; width="1em"></mspace>\mathrm{M}<mspace\; width="1em"></mspace>{\mathrm{Ca}}^{2+}+{10}^{-4}<mspace\; width="1em"></mspace>{\mathrm{MF}}^{-}$

4. ${10}^{-5}<mspace\; width="1em"></mspace>\mathrm{M}<mspace\; width="1em"></mspace>{\mathrm{Ca}}^{2+}+{10}^{-3}<mspace\; width="1em"></mspace>{\mathrm{MF}}^{-}$

Find the conditions that give the maximum yield of SO₃ for the exothermic reaction:

2SO₂(g) + O₂(g) ⇌ 2SO₃(g)

If a solution of 0.1 M NH₄OH and 0.1 M NH₄Cl has pH 9.25, then the pK_b of NH₄OH will be:

The compound with the highest pH among the following is:

1. CH₃COOK

2. Na₂CO₃ 

3. NH₄Cl

4. NaNO₃

The sharp pH change near the equivalence point in acid-base titration enables indicator detection. Which of the following equations correctly explains the pH change based on the concentration ratio of an indicator's conjugate acid (HIn) and base (In⁻) forms?

Find the solubility product (Kₛₚ) (in mol³ litre^-3) of M₂S having a molar solubility of 3.5 × 10⁻⁶ mol L⁻¹.

1. $1.7\times {10}^{-6}$

2. $1.7\times {10}^{-16}$

3. $1.7\times {10}^{-18}$

4. $1.7\times {10}^{-12}$

Find K_sp of an MX₂ type electrolyte if its solubility is 0.5 × 10⁻⁴ mol L⁻¹.

1. $5\times {10}^{-12}$

2. $25\times {10}^{-10}$

3. $1\times {10}^{-13}$

4. $5\times {10}^{-13}$