Equilibrium practice · Medium

Statement-I: $\mathrm{P}^{\mathrm{H}}$ of a solution changes from 6 to 7 when diluted by 10 times

Statement-II: If $\left[\mathrm{H}^{+}\right]$decreases 10 times, $\mathrm{P}^{\mathrm{H}}$ increases by one unit .

The equilibrium, $\mathrm{N}_2(\mathrm{~g})+3 \mathrm{H}_2(\mathrm{~g}) \rightleftharpoons 2 \mathrm{NH}_3(\mathrm{~g})$ is attained at $25^{\circ} \mathrm{C}$ in a closed container and an inert gas $\mathrm{He}$ is introduced. Choose the correct statement.

In an aqueous solution of volume $500 \mathrm{ml}$, when the reaction of $2 \mathrm{Ag} \mathrm{Cu}=\mathrm{Cu}^{2+}+2 \mathrm{Ag}$ reached equilibrium the $\left[\mathrm{Cu}^{2+}\right]$ was $\mathrm{x}$ M. When $500 \mathrm{ml}$ of water is further added, at the equilibrium $\left[\mathrm{Cu}^{2+}\right]$ will be

A buffer solution contains 0.1 mole of sodium acetate dissolved in $1000\,\mathrm{c}{\mathrm{m}^3}$ of 0.1 M acetic acid. To the above buffer solution, 0.1 mole of sodium acetate is further added and dissolved. The pH of the resulting buffer is

$\mathrm{p}{K_a}$ of a weak acid is defined as

Assertion :- pH of a buffer changes with temperature.

Reason :- Ionic product of water $(K_w)$ changes with temperature.

Calculate ionic product of ${H}_2 {O}$ at $50^{\circ} \mathrm{C}$.

Aqueous solution is alkaline in case of

$\mathrm{pH}$ of $0.01 \text{ M } \left(\mathrm{NH}_4\right)_2 \mathrm{SO}_4$ and $0.02 \mathrm{~M} ~\mathrm{NH}_4 \mathrm{OH}$ buffer $\left(\text{pK}_{\mathrm{a}}\left[\mathrm{NH}_4{ }^{+}\right]=9.26\right)$ is

$2 {KClO}_3(\mathrm{s}) \rightleftharpoons 2 {KCl}(\mathrm{s})+3 {O}_2(\mathrm{g})$

In the equilibrium given above, the equilibrium constant for the reverse reaction increases as the temperature of the system increases. Which is the correct statement about the system:-

Statement (S): $\mathrm{Cu}^{+}, \mathrm{Ag}^{+}$are noft acids whereas $\mathrm{CN}^{-}, \mathrm{H}^{-}$are soft bases.

Explanation (E): Soft acids do not possess noble configuration whereas soft bases have donor atom of easily polarised nature.

For the reaction, $\mathrm{PC}{\mathrm{l}_3}\left( \mathrm{g} \right) + \mathrm{C}{\mathrm{l}_2}\left( \mathrm{g} \right)\;\leftrightharpoons\;\mathrm{PC}{\mathrm{l}_5}\left( \mathrm{g} \right)$ the value of ${K_c}$ at $250^\circ C$ is 26. The value of ${K_p}$ at this temperature will be

Statement I: Aqueous solutions of ${NaCl}$ and ${CH}_3 {COONH}_4$ are neutral $\left(\mathrm{pH}=7.0\right.$ at $25^{\circ} \mathrm{C}).$

Statement II: Both the salts do not undergo hydrolysis.

STATEMENT-1 : Sparingly soluble salts $A B$ and $X Y_2$ with the same solubility product, will have different solubility.

STATEMENT-2 : Solubility of sparingly soluble salt depend upon solubility product.

STATEMENT I : Solubility of sparingly soluble salt decreases due to common ion effect.

STATEMENT II : Solubility product concentare does not depend on common ion effect.

Statement I: If equilibrium constant is $\mathrm{N}_2+\mathrm{O}_2 \rightleftharpoons 2 \mathrm{NO}$ is $\mathrm{K}$, the equilibrium constant for $\frac{1}{2} N_2+\frac{1}{2} O_2 \rightleftharpoons N O$ is $\mathrm{K} 2$

Statement II : Equilibrium constant does not depend on the stoichiometry form of equation

Solubility product of the salt, $A_x B_y$ will be represented most suitably, if the solubility is represented by $S$

Assertion (A): If the equation for a reaction is reversed, the equilibrium constant is inverted and if the equation is multiplied by 2, the equilibrium constant is squared.

Reason (R): The numerical value of an equilibrium constant depends on the way the equation for the reaction is written.

Statement-I: The $p^N$ of a buffer solution containing equal moles of acetic ncid and sodium ncetate is 4.8

$\text { ( } P^{k_a} \text { of of acetic acid is } 4.8 \text { ) }$

Statement-II: The ionic product of water at $25^{\circ} \mathrm{C}$ is $10^{-14} \mathrm{~mol}^2 \mathrm{lit}^{-2}$

If $10\,ml$ of 0.4 m $C{H_3}COOH$ solution is mixed with $40\,mol$ of 0.1 m $NaOH$ solution then the ${p^H}$ of the resultant solution becomes (given ${p^{ka}}\,of\,C{H_3}COOH = 4.75$)

For the reaction, ${N_2}\;\; + \;3{H_2}\;\leftrightharpoons\;2N{H_3};$ in a vessel, after the addition of equal number of mole of ${N_2}$ and ${H_2}$, equilibrium state is formed. Which of the following is correct?

Given ${K_{SP}}\left( {Ag\,I} \right) = 8.5 \times {10^{ - 17}}$. The solubility of $AgI$ in $0.1 M$ $KI$ solution is

Statement-I: In the reaction, $I_2+\mathrm{I}^{-} \rightarrow \mathrm{I}_3^{-}, \mathrm{I}_2$ acts as Lewis base.

Statement-II: In this reaction I' donates an electron pair for sharing with iodine

The following reaction was allowed to come to equilibrium: $\mathrm{A}(\mathrm{g})+2 \mathrm{~B}(\mathrm{g}) \rightleftharpoons \mathrm{C}(\mathrm{g}).$ The initial amounts of reactants placed into a $5 \mathrm{~L}$ in both vessels were $1 \mathrm{~mol}$ of $\mathrm{A}$ and $1.8 \mathrm{~mol}$ of $\mathrm{B}.$ After the reaction reached equilibrium, $1 \mathrm{~mol}$ of $\text{B}$ was found. Calculate the value of $K_C$ for this reaction.

Read the following statements carefully and choose the correct answer

(i) Water and water vapour remain in equilibrium position at atmospheric pressure ( $1.013 \,{bar})$ and at $100^{\circ} {C}$ in a closed vessel.

(ii) The boiling point of water is $100^{\circ} {C}$ at 1.013 bar pressure

(iii) Boiling point of the liquid depends on the atmospheric pressure.

(iv) Boiling point depends on the altitude of the place; at high altitude the boiling point increases.

Assertion :- The precipitate the cations of fourth group in qualitative analysis, medium is made alkaline before passing $\mathrm{H}_2 \mathrm{S}$ gas.

Reason :- This is done to suppress the ionization of $\mathrm{H}_2 \mathrm{S}.$

Which of the following solutions cannot act as a buffer?

The $\left[\mathrm{H}^{+}\right]$of a resulting solution that is $0.01 \mathrm{M}$ acetic acid $\left(K_a=1.8 \times 10^{-5}\right)$ and $0.01 \mathrm{M}$ in benzoic acid $\left(K_a=6.3 \times 10^{-5}\right)$ :

$\mathrm{A}_2 \mathrm{B}$ dissociates as $\mathrm{A}_2 \mathrm{B}(\mathrm{s}) \rightleftharpoons 2 \mathrm{~A}(\mathrm{g})+\mathrm{B}(\mathrm{g})$.

At equilibrium total pressure is $P$, then $K_p$ for the reaction is

(A): The addition of silver ions to a mixture of aqueous sodium chloride and sodium bromide solution will first precipitate ${AgBr}$ rather than ${AgCl}$

(R): The value of $K_{s p}$ of ${AgCl} < K_{s p}$ of ${AgBr}$