Equilibrium Questions

A buffer solution is prepared in which the concentration of ${\text{NH}}_3$ is $0.30\text{M}$ and the concentration of ${\text{NH}}_4^{+}$ is $0.20\text{M}$. If the equilibrium constant, $K_b$ for ${\text{NH}}_3$ equals $1.8\times 10^{-5}$, what is the pH of this solution?

For the reaction ${\text{N}}_2(\text{g})+{\text{O}}_2(\text{g})\rightleftharpoons 2\text{NO}(\text{g})$, the equilibrium constant is $K_1$. The equilibrium constant is $K_2$ for the reaction $2\text{NO}(\text{g})+{\text{O}}_2(\text{g})\rightleftharpoons 2{\text{NO}}_2(\text{g})$. What is $K$ for the reaction ${\text{NO}}_2(\text{g})\rightleftharpoons \frac{1}{2}{\text{N}}_2(\text{g})+{\text{O}}_2(\text{g})$?

If the $E_{cell}^{\circ }$ for a given reaction has a negative value, then which of the following gives the correct relationships for the values of $\Delta G^{\circ }$ and $K_{eq}$?

Which of the following is least likely to behave as a Lewis base?

The value of $\Delta H$ for the reaction ${\text{X}}_{2(\text{g})}+4{\text{Y}}_{2(\text{g})}\rightleftharpoons 2{\text{XY}}_{4(\text{g})}$ is less than zero. Formation of ${\text{XY}}_{4(\text{g})}$ will be favoured at:

For the reaction,${\text{CH}}_4(\text{g})+2{\text{O}}_2(\text{g})\rightleftharpoons {\text{CO}}_2(\text{g})+2{\text{H}}_2\text{O}(\text{l})$${\Delta }_{r}H=-170.8{\text{kJ mol}}^{-1}$ Which of the following statements is not true?

Which of the following pairs constitutes a buffer?

The hydrogen ion concentration of a $10^{-8}\text{M}$ $\text{HCl}$ aqueous solution at $298\text{K}$ ($K_w=10^{-14}$) is :

A weak acid, $\text{HA}$, has a $K_a$ of $1.00\times 10^{-5}$. If $0.100\text{ mole}$ of this acid is dissolved in one litre of water, the percentage of acid dissociated at equilibrium is closest to:

Standard entropies of $X_2,Y_2$ and $X{Y}_3$ are $60,40$ and $50{\text{JK}}^{-1}{\text{mol}}^{-1}$ respectively. For the reaction $\frac{1}{2}{X}_2+\frac{3}{2}{Y}_2\rightleftharpoons X{Y}_3,\Delta H=-30\text{kJ}$ to be at equilibrium, the temperature should be:

A solution contains ${\text{Ag}}^{+}$ and ${\text{Pb}}^{2+}$ at $0.10\text{M}$. Aqueous $\text{HCl}$ is added until $[{\text{Cl}}^{-}]=0.10\text{M}$. What are the concentrations of ${\text{Ag}}^{+}$ and ${\text{Pb}}^{2+}$ at equilibrium? ($K_{sp}$ for $\text{AgCl}=1.8\times 10^{-10}$, $K_{sp}$ for ${\text{PbCl}}_2=1.7\times 10^{-5}$)

Buffer solutions have constant acidity and alkalinity because:

Which of the following salts will give highest $pH$ in water?

For the reversible reaction: ${\text{N}}_2(g)+3{\text{H}}_2(g)\rightleftharpoons 2{\text{NH}}_3(g)+\text{heat}$, the equilibrium shifts in the forward direction:

For a given exothermic reaction, $K_P$ and $K_P^{\prime }$ are the equilibrium constants at temperatures $T_1$ and $T_2$, respectively. Assuming that the heat of reaction is constant in the temperature range between $T_1$ and $T_2$, and $T_2>T_1$, it is observed that: