Equilibrium practice · Mixed difficulty

In which of the following solvents, will $\mathrm{AgBr}$ have highest solubility?

pH of aqueous NaCI solution after electrolysis

$10 \mathrm{~mL}$ of $10^{-4} \mathrm{NH}_2 \mathrm{SO}_4$ is diluted with water to have 1 $\mathrm{dm}^3$ solution : The $\mathrm{pH}$ of this solution is :

Statement I : The hydrolysis of an ester in acidic medium does not change with pressure.

Statement II : Pressure does not show effect on equilibrium reaction taking place in solution

Statement (S): Acidic nature of boron trihalides is in the order :
$\mathrm{BF}_3<\mathrm{BCl}_3<\mathrm{BBr}_3<\mathrm{BI}_3$

Explanation (E): Basic nature of nitrogen trihalides is in the order :
$\mathrm{NF}_3>\mathrm{NCl}_3>\mathrm{NBr}_3>\mathrm{NI}_3$

STATEMENT I : $p H$ of aqueous solution of $\mathrm{NH}_4 \mathrm{Cl}$ is greater than 7.

STATEMENT II : Solution of a salt of weak base and strong acid is basic.

Indicate the correct answer out of the following for the reaction,

$\mathrm{N}{\mathrm{H}_4}\mathrm{Cl} + {\mathrm{H}_2}\mathrm{O}\;\leftrightharpoons\mathrm{N}{\mathrm{H}_4}\mathrm{OH} + \mathrm{HCl}$

Given ${K}_{\mathrm{a}}$ values are $5.76 \times 10^{-10}$ and $4.8 \times 10^{-10}$ for ${NH}_4^{+}$and ${HCN}$ respectively. What is the equilibrium constant for the following reaction ?

${NH}_4^{+}(aq.) +{CN}^{-}(aq.) \rightleftharpoons {NH}_3 (aq.) +{HCN} (aq.)$

Equal volumes of two solutions of a strong acid having ${P^H} = 3\,and\,{P^H} = 4$ are mixed together. The ${P^H}$ of the resulting solution is $\left[ {\log 0.5 = - 0.3} \right]$

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

In which of the following case, the reaction will proceed in backward direction:

$\mathrm{H}_2(\mathrm{g})+\mathrm{I}_2(\mathrm{g}) \rightleftharpoons 2 \mathrm{HI}(\mathrm{g}), \mathrm{K_c}=0.25 \text { at } 298 \mathrm{~K}$

Which one is not correct statement about buffer mixture?

(A): $\mathrm{BaCO}_3$ is more soluble in $\mathrm{HNO}_3$ than in plain water.

(R): Carbonate is a weak base and reacts with $\mathrm{H}^{+}$from the strong acid, causing the barium salt to dissociate.

An acid base indicator which is a weak acid has a $\mathrm{p} K_{\mathrm{a}}$ value $=5.5$. At what concentration ratio of sodium acetate to acetic acid would the indicator show a colour half way between those of its acid and conjugate base forms? The $\mathrm{p} K_{\mathrm{a}}$ value of acetic acid $=4.75.$
$[\operatorname{Antilog}(0.75)=5.62, \text{Antilog}(0.79)= 6.3, \operatorname{Antilog}(0.69)=4.93].$

$2.5\;mL\;of\;\frac{2}{5}\;M$ weak monoacidic base ${K_b} = 1 \times {10^{ - 12}}\;\;\mathrm{at}\;25^\circ C$ is titrated with $\frac{2}{{15}}\;M\;\mathrm{HCl}$ in water at $25^\circ C$ The concentration of ${\mathrm{H}^ + }$ at equivalence point is

${K_w} = 1 \times {10^{ - 14}}\;\mathrm{at}\;25^\circ C$

The solubility product of $\mathrm{A}{\mathrm{s}_2}{\mathrm{S}_3}\;\mathrm{is}\;2.8 \times {10^{ - 72}}$. What is the solubility of $\mathrm{A}{\mathrm{s}_2}{\mathrm{S}_3}$?

The vapour density of completely dissociated $\mathrm{N}{\mathrm{H}_4}\mathrm{Cl}$ would be:

Statement I : As per Lechutelier's principle, addition of heat to an equilibrium solid $\rightleftharpoons$ liquid results in decrease in the amount of solid.

Statement II : Reaction is endothermic, so on heating forward reaction is favoured.

For a chemical reaction ;
${A}(g)+{B}(\ell) \rightleftharpoons {D}(g)+{E}(g)$

Hypothetically at what temperature, ${K}_{{p}}={K}_{{c}}$ (when, ${R}=0.08 \ell-{atm} / {mole}-{K}$ )

Two buffers, $\mathrm{X}$ and $\mathrm{Y}$ of $\mathrm{pH}, 4.0$ and $6.0,$ respectively, are prepared from acid $\mathrm{HA}$ and the salt $\mathrm{NaA}.$ Both the buffers are $0.50 \mathrm{~M}$ in $\text{HA}.$ What would be the $\mathrm{pH}$ of the solution obtained by mixing equal volumes of the two buffers? The value of $K_{{a}}$ for $\mathrm{HA}=1.0 \times 10^{-5}(\log 5.05=0.7).$

A: Increasing the concentration of $\mathrm{H}_2$ will increase magnitude of equilibrium constant of the reaction $\mathrm{H}_2+\mathrm{I}_2 \rightleftharpoons 2 \mathrm{HI}$
R: Value of $\mathrm{K}_{\mathrm{C}}$ depends upon the concentration of reactants and products taken.

The $\mathrm{pH}$ of $0.005 \mathrm{~M}-\mathrm{NaOH}$ solution is $(\log 2=0.3)$

A mixture contains $0.1 \mathrm{~M} \mathrm{~HCOOH}\left(K_{\mathrm{a}}=1.8 \times 10^{-4}\right)$ and $0.1 \mathrm{~M} \mathrm{~HOCN}\left(K_{\mathrm{a}}=3.3 \times 10^{-4}\right)$. Thus, molar concentration of $\mathrm{H}_3 \mathrm{O}^{+}$is

In a reversible reaction $A+B \rightleftharpoons 2 C$, the equilibrium concentrations of (A), (B) and (C) are $2 \mathrm{M}, 2 \mathrm{M}$ and $4 \mathrm{M}$ respectively then $\mathrm{K}_{\mathrm{C}}$ for the reaction is

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

The $\mathrm{pH}$ of $\mathrm{0}\mathrm{.1}\;\mathrm{M}$ solution of the following salts increases in the order

Blue colour disappears in which solution by passing ${SO}_2$

Which acts both as Lowry Bronsted acid and base?

An example of a salt dissolved in water to give acidic solution is

Statement (S): HCl acts as weak base in liquid HF but weak acid in acetic acid
Explanation (E): The tendency to act as acid or base also depends upon the nature of other substance to accept or donate a proton.