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The dissociation constants for acetic acid and HCN at 25 °C are $1.5 \times 10^{-5}$ and $4.5 \times 10^{-10}$, respectively. The equilibrium constant for the equilibrium, $CN^- + CH_3COOH \rightleftharpoons HCN + CH_3COO^-$ would be:
The oxide that is not expected to react with sodium hydroxide is:
Which one of the following ionic species has the greatest proton affinity to form a stable compound?
The equilibrium constant $K_p$ for the following reaction is: $\text{MgCO}_3(s) \rightleftharpoons \text{MgO}(s) + \text{CO}_2(g)$
Which will make basic buffer?
The concentration of $Ag^+$ ions in a saturated solution of $Ag_2C_2O_4$ is $2.2 \times 10^{-4} \text{ mol L}^{-1}$. The solubility product of $Ag_2C_2O_4$ is:
Incorrect statement about pH and $[\text{H}^+]$ is:
The solubility product of a sparingly soluble salt $\text{AX}_2$ is $3.2 \times 10^{-11}$. Its solubility (in moles/litre) is:
The solubility of AgCl(s) with solubility product $1.6 \times 10^{-10}$ in 0.1 M NaCl solution would be?
For a reaction, $\text{BaO}_2(s) \rightleftharpoons \text{BaO}(s) + \text{O}_2(g)$; $\Delta H = +\text{ve}$. At equilibrium condition, the pressure of $\text{O}_2$ depends on the:
If a solution of $0.1 \text{ N NH}_4\text{OH}$ and $0.1 \text{ N NH}_4\text{Cl}$ has $\text{pH } 9.25$, then $pK_b$ of $\text{NH}_4\text{OH}$ is:
The $\text{p}K_b$ of dimethylamine and $\text{p}K_a$ of acetic acid are $3.27$ and $4.77$ respectively at $T\text{ (K)}$. The correct option for the pH of dimethylammonium acetate solution is:
Among the following, the correct order of acidity is:
The value of the equilibrium constant for a particular reaction is $1.6 \times 10^{12}$. When the system is in equilibrium, it will include:
Which of these is least likely to act as a Lewis base?
The equilibrium reaction that doesn't have equal values for $K_c$ and $K_p$ is:
Which of the following molecules acts as a Lewis acid?
The following equilibrium constants are given: $N_2 + 3H_2 \rightleftharpoons 2NH_3$; $K_1$ $N_2 + O_2 \rightleftharpoons 2NO$; $K_2$ $H_2 + \frac{1}{2}O_2 \rightleftharpoons H_2O$; $K_3$ The equilibrium constant for the oxidation of $NH_3$ by oxygen to give NO is:
The following pair constitutes a buffer is:
For the reaction $3\text{O}_2(g) \rightleftharpoons 2\text{O}_3(g)$ at $298\text{ K}$, $K_c$ is found to be $3.0 \times 10^{-59}$. If the concentration of $\text{O}_2$ at equilibrium is $0.040\text{ M}$, then the concentration of $\text{O}_3$ in $\text{M}$ is: