If the equilibrium constant for N₂(g) + O₂ (g) ⇄ 2NO(g) is K, the equilibrium constant for \(\frac{1}{2}\)N₂(g) + \(\frac{1}{2}\)O₂(g) ⇄ NO(g) will be?

1. $K^{\frac{1}{2}}$

2. $\frac{1}{2}K$

3. K

4. $K^2$

The value of the equilibrium constant for a particular reaction is 1.6 × 10¹². When the system is in equilibrium, it will include:
1. All reactants
2. Mostly reactants
3. Mostly products
4. Similar amounts of reactants and products

The K_sp of Ag₂CrO₄, AgCl, AgBr, and Agl are respectively, 1.1 × 10^–12, 1.8 × 10^–10, 5.0 × 10^–13, 8.3 × 10^–17. Which one of the following salts will precipitate last if ${\mathrm{AgNO}}_3$ solution is added to the solution containing equal moles of NaCl, NaBr, Nal, and Na₂CrO₄?

1. Agl

2. AgCl

3. AgBr

4. Ag₂CrO₄

For the reversible reaction:
N₂(g) + 3H₂(g) \(\rightleftharpoons\) 2NH₃(g) + heat 

The equilibrium shifts in a forward direction:

1. by increasing the concentration of $N{H}_3\left(g\right)$
2. by decreasing the pressure.
3. by decreasing the concentration of $N_2\left(g\right)$ $and$ $H_2\left(g\right)$
4. by increasing pressure and decreasing temperature.

Buffer solutions have constant acidity and alkalinity because:
 

A buffer solution is prepared in which the concentration of NH₃ is 0.30 M and the concentration of  $N{H}_4^{+\hspace{0.17em}}$ 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

What conditions favor the formation of 2XY₄(g) in the reaction X₂(g) + 4Y₂(g) ⇋ 2XY₄(g), considering that the enthalpy change (ΔH) is negative?

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

Calculate the hydrogen ion concentration, [\(\text{H}^+\)] (in \(\text{mol}/\text{L}\)), of a buffer solution prepared by mixing \(0.10 \text{ M}\) acetic acid (\(\text{CH}_3\text{COOH}\)) and \(0.20 \text{ M}\) sodium acetate (\(\text{CH}_3\text{COONa}\)), given that the acid dissociation constant (\(\text{K}_a\)) for acetic acid is \(1.8 \times 10^{-5}\):

1. \(3 . 5 \times 10^{- 4}\)
2. \(1 . 1 \times 10^{- 5}\)
3. \(1 . 8 \times 10^{- 5}\)
4. \(9 . 0 \times10^{- 6}\)${\mathrm{}}^{}$

The equilibrium reaction that doesn't have equal values for K_c and K_p is: 

1. \(2NO(g) \rightleftharpoons N_2(g) + O_2(g)\)
2. \(SO_2(g) + NO_2(g) \rightleftharpoons SO_3(g) + NO(g)\)
3. \(H_2(g) + I_2(g) \rightleftharpoons 2HI (g)\)
4. \(2C(s) + O_2(g) \rightleftharpoons 2CO_2(g)\)