When equal volumes of the following solutions are mixed, in which case will AgCl (K_sp = 1.8×10^–10) precipitate?

1. 10^–4 M Ag⁺ and 10^–4 M Cl^–

2. 10^–5 M Ag⁺ and 10^–5 M Cl^–

3. 10^–6 M Ag⁺ and 10^–6 M Cl^–

4. 10^–10 M Ag⁺ and 10^–10 M Cl^–

If a weak base has the dissociation constant, $K_b$, then the value of the dissociation constant, $K_a$, of its conjugate acid is given by:

1. $1/K_{b }$

2. $K_w/K_{b }$

3. $K_b/K_{w }$

4. $1_{}/K_{w }$

An aqueous solution of NaOH has a concentration of 0.01 mol/L.
Calculate the pH of the NaOH solution at 25 °C.

Given the ionic product of water is K_w = [H⁺] [OH^–] = 10^–14 mol²/L² (at 25 °C)
1.  11

2.  7

3.  14

4.  12

Which of the following reactions satisfies the condition \(K_p>K_c \text { ? }\)

1. H₂(g) + I₂(g) → 2HI(g)

2. N₂(g) + 3H₂(g) → 2NH₃(g)

3. 2SO₃(g) → 2SO₂(g) + O₂(g)

4. PCl₃(g) + Cl₂(g) → PCl₅(g)

For the reaction, A(g) + 2B(g) ⇌ 3C, K_c = 10². If at equilibrium one mole of A is added and 1 mole of C is removed then, the value of K_c will be: 

Which of the following species can act as both Bronsted acid and Bronsted base in an aqueous solution?

pH of a saturated solution of $\mathrm{Ca}{\left(\mathrm{OH}\right)}_2$ is 9. The solubility product $\left({\mathrm{K}}_{\mathrm{sp}}\right)$ of $\mathrm{Ca}{\left(\mathrm{OH}\right)}_2$ is:

1. $0.5\times {10}^{-10}$

2. $0.5\times {10}^{-15}$

3. $0.25\times {10}^{-10}$

4. $0.125\times {10}^{-15}$

For the reaction ${\mathrm{N}}_2\left(\mathrm{g}\right)$ $+$ ${\mathrm{O}}_2\left(\mathrm{g}\right)\leftrightharpoons 2\mathrm{NO}\left(\mathrm{g}\right)$ the equilibrium constant is K₁. The equilibrium constant is K₂ for the reaction  $2\mathrm{NO}\left(\mathrm{g}\right)$ $+$ ${\mathrm{O}}_2\left(\mathrm{g}\right)$ $\leftrightharpoons 2{\mathrm{NO}}_2\left(\mathrm{g}\right)$ $$
The value of K for the reaction given below will be:
${\mathrm{NO}}_2\left(\mathrm{g}\right)\leftrightharpoons \frac{1}{2}{\mathrm{N}}_2\left(\mathrm{g}\right)$ $+{\mathrm{O}}_2\left(\mathrm{g}\right)$ $$

1.  $\frac{1}{4}$ $\left(4\right)$ ${\mathrm{K}}_1$ ${\mathrm{K}}_2$

2.  ${\left[\frac{1}{{\mathrm{K}}_1{\mathrm{K}}_2}\right]}^{1/2}$

3.  $\frac{1}{\left({\mathrm{K}}_1{\mathrm{K}}_2\right)}$

4.  $\frac{1}{\left(2{\mathrm{K}}_1{\mathrm{K}}_2\right)}$

A buffer solution is prepared in which the concentration of NH₃ is 0.30 M and the concentration of  $N{H}_4^{+ }$ is 0.20 M. If the equilibrium constant, K_b for NH₃ equals 1.8×10^–5, then what is the pH of this solution? 
(log 1.8 = 0.25; log 0.67 = –0.176)

1.  9.43
2.  11.72
3.  8.73
4.  9.08