Equilibrium practice · Mixed difficulty

STATEMENT I : $p H$ of $10 \mathrm{M}$ $HCl$ aqucous solution is less than 1.

STATEMENT II : $p H$ is equal to logarithm of $\mathrm{H}^{+}$concentration.

The exothermic formation of $\mathrm{Cl}{\mathrm{F}_3}$ is represented by the equation
$\mathrm{C}{\mathrm{l}_2}\left( \mathrm{g} \right) + 3{\mathrm{F}_2}\left( \mathrm{g} \right)\leftrightharpoons2\mathrm{Cl}{\mathrm{F}_3}\left( \mathrm{g} \right);$
$\Delta H = - 329\,\mathrm{kJ}$
Which of the following will increase the quantity of $\mathrm{Cl}{\mathrm{F}_3}$ in an equilibrium mixture of $\mathrm{C}{\mathrm{l}_2},{\mathrm{F}_2}\mathrm{and Cl}{\mathrm{F}_3}$?

Hydrogen ion concentration in mol/L in a solution of $\mathrm{pH} = 5.4$ will be

$0.15$ mole of CO taken in a $2.5 \mathrm{~L}$ flask is maintained at $750 \mathrm{~K}$ along with a catalyst so that the following reaction can take place.

$\mathrm{CO}(g)+2 \mathrm{H}_2(g) \rightleftharpoons \mathrm{CH}_3 \mathrm{OH}(g)$

Hydrogen is introduced until the total pressure of the system is $8.5 \mathrm{~atm}$ at equilibrium and 0.08 mole of $\mathrm{CH}_3 \mathrm{OH}$ are formed. The value of $K_c$ is ..... $\mathrm{mol}^{-2} \mathrm{~L}^{-2}$ (Assigning ideal behaviour).

$A B_3(g)$ is dissociates as $A B_3(g) \rightleftharpoons A B_2(g)+\frac{1}{2} B_2(g)$.
When the initial pressure of $A B_2$ is 800 torr and the total pressure developed at equilibrium is 900 torr. What fraction of $A B_3({~g})$ is dissociated?

The ionisation of strong electrolytes in acetic acid, compared to in water, is

For the following Assertion and Reason the correct option is
Assertion: The pH of water increases with increase in temperature.
Reason: The dissociation of water into ${H^ + }$ and $O{H^ - }$ is an exothermic reaction.

If the solubility product of lead iodide $\left( {\mathrm{Pb}{\mathrm{I}_2}} \right)$ is $3.2 \times {10^{ - 8}},$ its solubility will be:

If $K_1$ and $K_2$ are the equilibrium constants of the equilibria (I) and (II) respectively, what is the relationship between the two constants?

(I) $\mathrm{SO}_2(\mathrm{g})+\frac{1}{2} \mathrm{O}_2(\mathrm{g}) \rightleftharpoons \mathrm{SO}_3(\mathrm{g}) ; K_1$

(II) $2 \mathrm{SO}_3(\mathrm{g}) \rightleftharpoons 2 \mathrm{SO}_2(\mathrm{g})+\mathrm{O}_2(\mathrm{g}) ; K_2$

Statement-I: pH of a buffer solution does not change on dilution.

Statement-II:On dilution the ratio of concentration of salt and acid (or base) remains unchanged.

${K}_{{P}}$ and ${K}_{{C}}$ ratio for the following reaction

${SO}_2({~g})+\frac{1}{2} {O}_2({~g}) \rightleftharpoons {SO}_3({~g})$ will be

Statement I In the titration of weak acid with strong base,
neutral point can't be attained
Statement II Salt formed is hydrolysed to form alkaline
solution

$20 \mathrm{~mL}$ of $0.1 \mathrm\,\,{M} \mathrm\,\,\,{NH}_4 \mathrm{OH}$ is mixed with $40 \mathrm{~mL}$ of $0.05 \mathrm\,\,{M} \mathrm\,\,\,{HCl}$. The $\mathrm{pH}$ of the mixture is nearest to
(Given : $\mathrm{K}_{\mathrm{b}}\left(\mathrm{NH}_4 \mathrm{OH}\right)=1 \times 10^{-5}, \log 2=0.30, \log 3=$ $0.48, \log 5=0.69, \log 7=0.84, \log 11=1.04$ )

In the dissociation of $2\mathrm{HI}\leftrightharpoons{\mathrm{H}_2} + {\mathrm{I}_2},$ the degree of dissociation will be influenced by the:

In ${A}+{B} \rightleftharpoons {C}$. The unit of equilibrium constant is :

STATEMENT I : Addition of $\mathrm{HCl}(a q)$ to $\mathrm{HCOOH}(a q)$ decreases the ionization of $\mathrm{HCOOH}$

STATEMENT II : Common ion effect of $\mathrm{H}^{+}$ion, reduces the ionization of $\mathrm{HCOOH}$

At dissociation constant of heavy water is $4 \times 10^{-15}$ at $35^{\circ} \mathrm{C}.$ If its density is $1.04 \mathrm{~g} / \mathrm{mL}.$ Calculate its ionic product & degree of dissociation.

Which of the following combinations of $\mathrm{CH}_3 \mathrm{COOH}+\mathrm{NaOH}$ would result in the formation of acidic buffer?

According to Le-chatlieris principle, if heat is given to solid-liquid system then

Assertion (A): On increasing temperature pH of $\mathrm{H}_2 \mathrm{O}$ decreases.

Reason (R): At high temperature water become acidic.

The equilibrium constant, $K$ for the reaction

$2\mathrm{HI}\left( \mathrm{g} \right)\;\leftrightharpoons\;{\mathrm{H}_2}\left( \mathrm{g} \right) + {\mathrm{I}_2}\left( \mathrm{g} \right)$

At room temperature is 2.85 and that at $698\;K$, it is $1.4 \times {10^{ - 2}}$. This implies that

Calculate the $\mathrm{pH}$ of solution made by two solution having $\mathrm{pH}=1.5$ and 2.5

The equilibrium reaction ${PCl}_5(s) \rightleftharpoons {PCl}_3(s)+{Cl}_2({g})$ is carried in a closed vessel. What is the effect on equilibrium concentration of ${Cl}_2({g})$ by adding ${PCl}_5$ at constant temperature?

$\mathrm{C}(\mathrm{s})+\mathrm{H}_2 \mathrm{O}(\mathrm{g}) \rightleftharpoons \mathrm{CO}(\mathrm{g})+\mathrm{H}_2(\mathrm{g}) ; \Delta \mathrm{H} < 0$. The above equilibrium will proceed in forward direction when

A solution that has reserve acidity (or) alkalinity is called

A volume of $10 \mathrm{~ml}$ of $0.1 \mathrm{~M}$ tribasic acid, $\mathrm{H}_3 \mathrm{A}$ is titrated with $0.1 \mathrm{~M}-\mathrm{NaOH}$ solution. What is the ratio (approximate value) of $\frac{\left[\mathrm{H}_3 \mathrm{A}\right]}{\left[\mathrm{A}^{3-}\right]}$ at the second equivalent point? Given: $K_1=7.5 \times 10^{-4} ; K_2=10^{-8} ; K_3=10^{-12}$

At $90^{\circ} \mathrm{C}$ pure water has $\left[\mathrm{H}_3 \mathrm{O}^{+}\right]=10^{-6} \mathrm{M}$. What is the value of $\mathrm{K}_{\mathrm{w}}$ at this temperature?

The melting points of most of the solid substances increase with an increase of pressure. However, ice melts at a temperature lower than its usual melting point when the pressure is increased. This is because:

The $pH$ of an aqueous solution of $\mathrm{C}{\mathrm{H}_3}\mathrm{COONa}$ of concentration $C\left( \mathrm{M} \right)$ is given by

Equal volume of two solutions having pH = 2, pH = 4 are mixed. The pH of the resulting solution is