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The concentration of the $Ag^+$ ions in a saturated solution of $Ag_2C_2O_4$ is $2.2 \times 10^{-4}\text{ M}$. The solubility product of $Ag_2C_2O_4$ is:

The percentage of pyridine ($C_5H_5N$) that forms pyridinium ion ($C_5H_5NH^+$) in a 0.10 M aqueous pyridine solution ($K_b$ for $C_5H_5N = 1.7 \times 10^{-9}$) is:

The $K_{sp}$ of $Ag_2CrO_4$, $AgCl$, $AgBr$ and $AgI$ are, respectively, $1.1 \times 10^{-12}$, $1.8 \times 10^{-10}$, $5.0 \times 10^{-13}$, and $8.3 \times 10^{-17}$. The salt that precipitates last if $AgNO_3$ solution is added to the solution containing equal moles of $NaCl$, $NaBr$, $NaI$ and $Na_2CrO_4$ is:

Which of the following statements is correct for a reversible process in a state of equilibrium?

The molecule that is least likely to behave as Lewis base is:

For the reversible reaction, $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g) + \text{Heat}$, the equilibrium shifts in the forward direction:

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 value of equilibrium constant of the reaction $HI(g) \rightleftharpoons \frac{1}{2} H_2(g) + \frac{1}{2} I_2(g)$ is $8.0$. The equilibrium constant of the reaction $H_2(g) + I_2(g) \rightleftharpoons 2HI(g)$ will be:

The oxide that is not expected to react with sodium hydroxide is:

What is the molarity of the saturated solution if the solubility product for a salt of type AB is $4 \times 10^{-8}$?

Which one of the following ionic species has the greatest proton affinity to form a stable compound?

Given that the equilibrium constant for the reaction $2\text{SO}_2(g) + \text{O}_2(g) \rightleftharpoons 2\text{SO}_3(g)$ has a value of $278$ at a particular temperature, the value of the equilibrium constant for the following reaction at the same temperature will be: $\text{SO}_3(g) \rightleftharpoons \text{SO}_2(g) + \frac{1}{2} \text{O}_2(g)$

Consider the following reaction: $\text{A}_2(g) + \text{B}_2(g) \rightleftharpoons 2\text{AB}(g)$. At equilibrium, the concentrations of $[\text{A}_2] = 3.0 \times 10^{–3} \text{ M}$; $[\text{B}_2] = 4.2 \times 10^{–3} \text{ M}$ and $[\text{AB}] = 2.8 \times 10^{–3} \text{ M}$. The value of $K_c$ for the above-given reaction in a sealed container at $527^\circ\text{C}$ is:

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?

In qualitative analysis, the metals of Group I can be separated from other ions by precipitating them as chloride salts. A solution initially contains $\text{Ag}^+$ and $\text{Pb}^{2+}$ at a concentration of $0.10 \text{ M}$. Aqueous $\text{HCl}$ is added to this solution until the $\text{Cl}^-$ concentration is $0.10 \text{ M}$. What will the concentration of $\text{Ag}^+$ and $\text{Pb}^{2+}$ at equilibrium? ($K_{sp}$ for $\text{AgCl} = 1.8 \times 10^{-10}$, $K_{sp}$ for $\text{PbCl}_2 = 1.7 \times 10^{-5}$)

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:

Acidified $K_2Cr_2O_7$ solution turns green when $Na_2SO_3$ is added to it. This is due to the formation of: