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?

1.	\(\log \left[\frac{\mathrm{In}^{-}}{\mathrm{HIn}}\right]=\mathrm{pK}_{\mathrm{In}}-\mathrm{pH} \)
2.	\(\log \left[\frac{\mathrm{HIn}}{\mathrm{In}}\right]=\mathrm{pK}_{\mathrm{In}}+\mathrm{pH} \)
3.	\(\log \left[\frac{\mathrm{HIn}}{\mathrm{In}}\right]=\mathrm{pH}-\mathrm{pK}_{\mathrm{In}} \)
4.	\(\log \left[\frac{\mathrm{In}}{\mathrm{HIn}}\right]=\mathrm{pH}-\mathrm{pK}_{\mathrm{In}}\)

Solubility of a M₂S salt is $3.5\times {10}^{-6}$ then find out solubility product.

1. $1.7\times {10}^{-6}$

2. $1.7\times {10}^{-16}$

3. $1.7\times {10}^{-18}$

4. $1.7\times {10}^{-12}$

The solubility of MX₂ type electrolyte is $0.5\times {10}^{-4}$ mol L^-1. Then find out K_sp of electrolytes:

1. $5\times {10}^{-12}$

2. $25\times {10}^{-10}$

3. $1\times {10}^{-13}$

4. $5\times {10}^{-13}$

If the solubility product of a sparingly soluble salt AX₂ is $3.2\times {10}^{-11}$ , find its solubility (in mol L^–1):

What is the degree of ionization of an acid solution with a concentration of 0.005 M and a pH of 5?
 

1. $0.1\times {10}^{-2}$

2. $0.2\times {10}^{-2}$

3. $0.5\times {10}^{-4}$

4. $0.6\times {10}^{-6}$

Which among the following acids is the strongest, based on their given pKa values?

(I) 4.0

(II) 3.5

(III) 2.5

(IV) 2.0

1. I.0

2. II

3. III

4. IV

A solution has hydrogen ion concentration 0.0005 M, its pOH is:

1. 8.2798

2. 10.6990

3. 12.7854

4. 13.3344

At 25°C, the pH of a solution containing 0.10 M sodium acetate and 0.03 M acetic acid will be:

[pK_a value of ${\mathrm{CH}}_3\mathrm{COOH}=4.57$]

1. 3.24

2. 4.59

3. 5.09

4. 6.67

Ionization constant of acetic acid(\(\mathrm{CH}_3 \mathrm{COOH}\)) is $1.8\times {10}^{-5}$. The concentration of H⁺ ions in 0.1 M solution is:

1. $1.8\times {10}^{-3}<mspace\; width="1em"></mspace>\mathrm{M}$

2. $2.8\times {10}^{-3}<mspace\; width="1em"></mspace>\mathrm{M}$

3. $1.3\times {10}^{-3}<mspace\; width="1em"></mspace>\mathrm{M}$

4. $1.34\times {10}^{-4}<mspace\; width="1em"></mspace>\mathrm{M}$

What is the ionization percentage for a weak acid with a dissociation constant of \(4.9×10^{−8}\) at 0.1 M concentration?