Equilibrium

Which of the following mixture gives a buffer solution with pH = 9.25 ?

Given : p$K_\mathrm{b}$ ($\ce{NH4OH}$) = 4.75

Consider the general reaction given below at 400 K

$\ce{xA(g) <=> yB(g)}$

The values of $K_\mathrm{p}$ and $K_\mathrm{c}$ are studied under the same condition of temperature but variation in x and y.

(i) $K_\mathrm{p} = 85.87$ and $K_\mathrm{c} = 2.586$ appropriate units

(ii) $K_\mathrm{p} = 0.862$ and $K_\mathrm{c} = 28.62$ appropriate units

The value of x and y in (i) and (ii) respectively are:

For the following gas phase equilibrium reaction at constant temperature,

$\ce{NH3(g) <=> 1/2 N2(g) + 3/2 H2(g)}$

If the total pressure is $\sqrt{3}$ atm and the pressure equilibrium constant ($K_\mathrm{P}$) is 9 atm, then the degree of dissociation is given as $(x\times10^{-2})^{-1/2}$.

The value of x is (Nearest integer)

$\ce{X2(g) + Y2(g) <=> 2Z(g)}$

$\ce{X2(g)}$ and $\ce{Y2(g)}$ are added to a 1 L flask and it is found that the system attains the above equilibrium at T(K) with the number of moles of $\ce{X2(g)}$, $\ce{Y2(g)}$ and Z(g) being 3, 3 and 9 mol respectively (equilibrium moles). Under these conditions of equilibrium, 10 mol of Z(g) is added to the flask and the temperature is maintained at T(K). Then the number of moles of Z(g) in the flask when the new equilibrium is established is . (Nearest integer).

Consider two Group IV metal ions $\ce{X^{2+}}$ and $\ce{Y^{2+}}$.

A solution containing 0.01 M $\ce{X^{2+}}$ and 0.01 M $\ce{Y^{2+}}$ is saturated with $\ce{H2S}$. The pH at which the metal sulphide YS will form as a precipitate is (Nearest integer)

(Given :

$K_\mathrm{sp}(\mathrm{XS}) = 1\times10^{-22}$ at 25$^\circ$C, $K_\mathrm{sp}(\mathrm{YS}) = 4\times10^{-16}$ at 25$^\circ$C,

[$\ce{H2S}$] = 0.1 M in solution, $K_\mathrm{a1}\times K_\mathrm{a2}$ ($\ce{H2S}$) $= 1.0\times10^{-21}$,

$\log 2 = 0.30$, $\log 3 = 0.48$, $\log 5 = 0.70$)

Consider the following gaseous equilibrium in a closed container of volume “V” at T(K).

$\ce{P2(g) + Q2(g) <=> 2PQ(g)}$

2 moles each of $\ce{P2(g)}$, $\ce{Q2(g)}$ and PQ (g) are present at equilibrium. Now one mole each of ‘$\ce{P2}$' and ‘$\ce{Q2}$' are added to the equilibrium keeping the temperature at T(K). The number of moles of $\ce{P2}$, $\ce{Q2}$ and PQ at the new equilibrium, respectively, are -

Consider a weak base ‘B' of $\mathrm{p}K_\mathrm{b} = 5.699$. ‘x' mL of 0.02 M HCl and ‘y' mL of 0.02 M weak base ‘B' are mixed to make 100 mL of a buffer of pH 9 at 25$^\circ$C. The values of ‘x' and ‘y' respectively are:

[Given: $\log 2 = 0.3010$, $\log 3 = 0.4771$, $\log 5 = 0.699$]

Consider the dissociation equilibrium of the following weak acid $\ce{HA <=> H^+(aq) + A^-(aq)}$

If the $\mathrm{p}K_\mathrm{a}$ of the acid is 4, then the pH of 10 mM HA solution is . (Nearest integer)

[Given : The degree of dissociation can be neglected with respect to unity]

The solubility product constants of $\ce{Ag2CrO4}$ and AgBr are 32x and 4y respectively at 298 K.

The value of $\left(\dfrac{{\text{molarity of }\ce{Ag2CrO4}}}{\text{molarity of AgBr}}\right)$ can be expressed as :

At 25$^\circ$C, 20.0 mL of 0.2 M weak monoprotic acid HX is titrated against 0.2 M NaOH. The pH of the solution (a) at the start of the titration (when NaOH has not been added) and (b) when 10 mL of NaOH is added respectively, are :

Given : $K_\mathrm{a} = 5\times10^{-4}$, p$K_\mathrm{a} = 3.3$, $\alpha \ll 1$

At T(K), the equilibrium constant of

$\ce{A2(g) + B2(g) <=> C(g)}$ is $2.7\times10^{-5}$.

What is the equilibrium constant for

$\ce{1/3 A2(g) + 1/3 B2(g) <=> 1/3 C(g)}$ at the same temperature?

The pH of a solution obtained by mixing 5 mL of 0.1 M $\ce{NH4OH}$ solution with 250 mL of 0.1 M $\ce{NH4Cl}$ solution is $\times\,10^{-2}$. (Nearest integer)

Given : p$K_\mathrm{b}$ ($\ce{NH4OH}$) $= 4.74$

$\log 2 = 0.30$

$\log 3 = 0.48$

$\log 5 = 0.70$

• (A)10 mL 0.2 M $\ce{Ca(OH)2}$ + 25 mL 0.1 M HCl
• (B)10 mL 0.01 M $\ce{H2SO4}$ + 10 mL 0.01 M $\ce{Ca(OH)2}$
• (C)10 mL 0.1 M $\ce{H2SO4}$ + 10 mL 0.1 M KOH

The values of pressure equilibrium constant recorded at different temperatures for the following equilibrium reaction have been given below:

$\ce{A(g) <=> B(g) + C(g)}$

The reaction $\ce{A(g) <=> B(g) + C(g)}$ was initiated with the amount ‘a' of A(g). At equilibrium it is found that the amount of A(g) remaining is $(a - x)$ at a total pressure of $p$.

The equilibrium constant $K_\mathrm{p}$ of the reaction can be calculated from the expression :

One mole each of He and A(g) are taken in a 10 L closed flask and heated to 400 K to establish the following equilibrium.

$\ce{A(g) <=> B(g)}$

$K_\mathrm{C}$ for this reaction at 400 K is 4.0. The partial pressures (in atm) of He and B(g) are respectively (at equilibrium).

(Assume He, A(g) and B(g) behave as ideal gases)

(Given : $R = 0.082$ L atm K$^{-1}$ mol$^{-1}$)

Consider the following reactions in which all the reactants and products are present in gaseous state

$2\mathrm{xy} \rightleftharpoons \mathrm{x_2} + \mathrm{y_2}$   $K_1 = 2.5\times10^{5}$

$\mathrm{xy} + \dfrac{1}{2}\mathrm{z_2} \rightleftharpoons \mathrm{xyz}$   $K_2 = 5\times10^{-3}$

The value of $K_3$ for the equilibrium

$\dfrac{1}{2}\mathrm{x_2} + \dfrac{1}{2}\mathrm{y_2} + \dfrac{1}{2}\mathrm{z_2} \rightleftharpoons \mathrm{xyz}$ is :

Solid carbon, CaO and $\ce{CaCO3}$ are mixed and allowed to attain equilibrium at T K.

$\ce{CaCO3 <=> CaO(s) + CO2(g)}$   Kp$_1 = 0.08$ atm

$\ce{C(s) + CO2(g) <=> 2CO(g)}$   Kp$_2 = 2$ atm

The partial pressure of CO is $\times 10^{-1}$ atm.