We know that the relationship between $K_c$ and $K_p$ is

$$K_p=K_c(R T)^{\Delta n}$$

What would be the value of $\Delta n$ for the reaction?

$$\mathrm{NH}_4 \mathrm{Cl}(\mathrm{s}) \rightleftharpoons \mathrm{NH}_3(\mathrm{~g})+\mathrm{HI}(\mathrm{g})$$

For the reaction, $\mathrm{H}_2(\mathrm{~g})+\mathrm{I}_2(\mathrm{~g}) \rightleftharpoons 2 \mathrm{HI}(\mathrm{g})$, the standard free energy is $\Delta G^{\ominus}>0$. The equilibrium constant $(K)$ would be

Which of the following is not a general characteristic of equilibria involving physical processes?

$\mathrm{PCl}_5, \mathrm{PCl}_3$, and $\mathrm{Cl}_2$ are at equilibrium at 500 K in a closed container and their concentrations are $0.8 \times 10^{-3} \mathrm{~mol} \mathrm{~L}^{-1}, 1.2 \times 10^{-3} \mathrm{~mol} \mathrm{~L}^{-1}$ and $1.2 \times 10^{-3} \mathrm{~mol} \mathrm{~L}^{-1}$, respectively. The value of $K_c$ for the reaction

$\mathrm{PCl}_5(\mathrm{~g}) \rightleftharpoons \mathrm{PCl}_3(\mathrm{~g})+\mathrm{Cl}_2(\mathrm{~g})$ will be

Which of the following statements is incorrect?

When hydrochloric acid is added to cobalt nitrate solution at room temperature, the following reaction takes place and the reaction mixture becomes blue. On cooling the mixture it becomes pink. On the basis of this information mark the correct answer.

$$\mathrm{\mathop {{{[Co({H_2}{O_6})]}^{3 + }}}\limits_{(Pink)} (aq) + 4C{l^ - }(aq)}\rightleftharpoons \mathrm{\mathop {{{[CoC{l_4}]}^{2 - }}}\limits_{(Blue)} (aq) + 6{H_2}O}(l)$$

The pH of neutral water at $25^{\circ} \mathrm{C}$ is 7.0 . As the temperature increases, ionisation of water increases, however, the concentration of $\mathrm{H}^{+}$ions and $\mathrm{OH}^{-}$ions are equal. What will be the pH of pure water at $60^{\circ} \mathrm{C}$ ?

The ionisation constant of an acid, $K_a$ is the measure of strength of an acid. The $K_a$ values of acetic acid, hypochlorous acid and formic acid are $1.74 \times 10^{-5}, 3.0 \times 10^{-8}$ and $1.8 \times 10^{-4}$ respectively. Which of the following orders of pH of $0.1 \mathrm{~mol} \mathrm{dm}^{-3}$ solutions of these acids is correct?

$K_{a_1}, K_{a_2}$ and $K_{a_3}$ are the respective ionisation constants for the following reactions.

$$\begin{aligned} & \mathrm{H}_2 \mathrm{~S} \rightleftharpoons \mathrm{H}^{+}+\mathrm{HS}^{-} \\ & \mathrm{HS}^{-} \rightleftharpoons \mathrm{H}^{+}+\mathrm{S}^{2-} \\ & \mathrm{H}_2 \mathrm{~S} \rightleftharpoons 2 \mathrm{H}^{+}+\mathrm{S}^{2-} \end{aligned}$$

The correct relationship between $K_{a_1}, K_{a_2}, K_{a_3}$ is

Acidity of $\mathrm{BF}_3$ can be explained on the basis of which of the following concepts?

Which of the following will produce a buffer solution when mixed i equal volumes?

In which of the following solvents is silver chloride most soluble?

What will be the value of pH of $0.01 \mathrm{~mol} \mathrm{dm}^{-3} \mathrm{CH}_3 \mathrm{COOH}$ $\left(K_a=1.74 \times 10^{-5}\right) ?$

$\mathrm{K}_{\mathrm{a}}$ for $\mathrm{CH}_3 \mathrm{COOH}$ is $1.8 \times 10^{-5}$ and $\mathrm{K}_{\mathrm{b}}$ for $\mathrm{NH}_4 \mathrm{OH}$ is $1.8 \times 10^{-5}$. The pH of ammonium acetate will be

Which of the following options will be correct for the stage of half completion of the reaction $\mathrm{A} \rightleftharpoons \mathrm{B}$ ?

On increasing the pressure, in which direction will the gas phase reaction proceed to re-establish equilibrium, is predicted by applying the Le-Chatelier's principle. Consider the reaction,

$$\mathrm{N}_2(\mathrm{~g})+3 \mathrm{H}_2(\mathrm{~g}) \rightleftharpoons 2 \mathrm{NH}_3(\mathrm{~g})$$

Which of the following is correct, if the total pressure at which the equilibrium is established, is increased without changing the temperature?

What will be the correct order of vapour pressure of water, acetone and ether at $30^{\circ} \mathrm{C}$ ? Given that among these compounds, water has maximum boiling point and ether has minimum boiling point?

At 500 K , equilibrium constant, $\mathrm{K}_{\mathrm{c}}$, for the following reaction is 5.

$$\frac{1}{2} \mathrm{H}_2(\mathrm{~g})+\frac{1}{2} \mathrm{I}_2(\mathrm{~g}) \rightleftharpoons \mathrm{HI}(\mathrm{g})$$

What would be the equilibrium constant $\mathrm{K}_{\mathrm{c}}$ for the reaction?

$$2 \mathrm{HI}(\mathrm{g}) \rightleftharpoons \mathrm{H}_2(\mathrm{~g})+\mathrm{I}_2(\mathrm{~g})$$

In which of the following reactions, the equilibrium remains unaffected on addition of small amount of argon at constant volume?

Assertion (A) Increasing order or acidity of hydrogen halides is $\mathrm{HF}<\mathrm{HCI}<\mathrm{HBr}<\mathrm{HI}$.

Reason (R) While comparing acids formed by the elements belonging to the same group of periodic table, H-A bond strength is a more important factor in determining acidity of an acid than the polar nature of the bond.

Assertion (A) A solution containing a mixture of acetic acid and sodium acetate maintains a constant value of pH on addition of small amounts of acid or alkali.

Reason (R) A solution containing a mixture of acetic acid and sodium acetate acts as buffer solution around pH 4.75 .

Assertion (A) The ionisation of hydrogen sulphide in water is low in the presence of hydrochloric acid.

Reason (R) Hydrogen sulphide is a weak acid.

Assertion (A) For any chemical reaction at a particular temperature, the equilibrium constant is fixed and is a characteristic property.

Reason (R) Equilibrium constant is independent of temperature.

Assertion (A) Aqueous solution of ammonium carbonate is basic.

Reason (R) Acidic/basic nature of a salt solution of a salt of weak acid and weak base depends on $\mathrm{K}_{\mathrm{a}}$ and $\mathrm{K}_{\mathrm{b}}$ value of the acid and the base forming it.

Assertion (A) An aqueous solution of ammonium acetate can act as a buffer.

Reason (R) Acetic acid is a weak acid and $\mathrm{NH}_4 \mathrm{OH}$ is a weak base.

Assertion (A) In the dissociation of $\mathrm{PCl}_5$ at constant pressure and temperature the addition of helium at equilibrium increases the dissociation of $\mathrm{PCl}_5$.

Reason (R) Helium removes $\mathrm{Cl}_2$ from the field of action.

For the reaction $\mathrm{N}_2 \mathrm{O}_4(\mathrm{~g}) \rightleftharpoons 2 \mathrm{NO}_2(\mathrm{~g})$, the value of K is 50 at 400 K and 1700 at 500 K . Which of the following option(s) is/are correct?

At a particular temperature and atmospheric pressure, the solid and liquid phases of a pure substance can exist in equilibrium. Which of the following term defines this temperature?

The ionisation of hydrochloric acid in water is given below

$$\mathrm{HCl}(\mathrm{aq})+\mathrm{H}_2 \mathrm{O}(l) \rightleftharpoons \mathrm{H}_3 \mathrm{O}^{+}(\mathrm{aq})+\mathrm{Cl}^{-}(\mathrm{aq})$$

Label two conjugate acid-base pairs in this ionisation.

The aqueous solution of sugar does not conduct electricity. However, when sodium chloride is added to water, it conducts electricity. How will you explain this statement on the basis of ionisation and how is it affected by concentration of sodium chloride?

$\mathrm{BF}_3$ does not have proton but still acts as an acid and reacts with $\ddot{\mathrm{N}} \mathrm{H}_3$. Why is it so? What type of bond is formed between the two?

Ionisation constant of a weak base MOH , is given expression

$$\mathrm{K}_{\mathrm{b}}=\frac{\left[\mathrm{M}^{+}\right]\left[\mathrm{OH}^{-}\right]}{[\mathrm{MOH}]}$$

Values of ionisation constant of some weak bases at a particular temperature are given below

$\begin{array}{ccccc}\text { Base } & \text { Dimethylamine } & \text { Urea } & \text { Pyridine } & \text { Ammonia } \\ \mathrm{K}_{\mathrm{b}} & 5.4 \times 10^{-4} & 1.3 \times 10^{-14} & \begin{array}{l}1.77 \times 10^{-9}\end{array} & 1.77 \times 10^{-5}\end{array}$

Arrange the bases in decreasing order of the extent of their ionisation at equilibrium. Which of the above base is the strongest?

Conjugate acid of a weak base is always stronger. What will be the decreasing order of basic strength of the following conjugate bases?

$$\mathrm{OH}^{-}, \mathrm{RO}^{-} \mathrm{CH}_3 \mathrm{COO}^{-}, \mathrm{Cl}^{-}$$

Arrange the following in increasing order of pH .

$$\left.\mathrm{KNO}_3(\mathrm{aq}), \mathrm{CH}_3 \mathrm{COONa}_{(\mathrm{aq}}\right) \mathrm{NH}_4 \mathrm{Cl}(\mathrm{aq}), \mathrm{C}_6 \mathrm{H}_5 \mathrm{COONH}_4(\mathrm{aq})$$

The value of $\mathrm{K}_{\mathrm{c}}$ for the reaction $2 \mathrm{HI}(\mathrm{g}) \rightleftharpoons \mathrm{H}_2(\mathrm{~g})+\mathrm{I}_2(\mathrm{~g})$ is $1 \times 10^{-4}$. At a given time, the composition of reaction mixture is $[\mathrm{HI}]=2 \times 10^{-5} \mathrm{~mol}$, $\left[\mathrm{H}_2\right]=1 \times 10^{-5} \mathrm{~mol}$ and $\left[\mathrm{I}_2\right]=1 \times 10^{-5} \mathrm{~mol}$. In which direction will the reaction proceed?

On the basis of the equation $\mathrm{pH}=-\log \left[\mathrm{H}^{+}\right]$, the pH of $10^{-8} \mathrm{~mol} \mathrm{dm}^{-3}$ solution of HCl should be 8 . However, it is observed to be less than 7.0. Explain the reason.

pH of a solution of a strong acid is 5.0 . What will be the pH of the solution obtained after diluting the given solution a 100 times?

A sparingly soluble salt gets precipitated only when the product of concentration of its ions in the solution ( $Q_{\mathrm{sp}}$ ) becomes greater than its solubility product. If the solubility of $\mathrm{BaSO}_4$ in water is $8 \times 10^{-4} \mathrm{~mol}$ $\mathrm{dm}^{-3}$. Calculate its solubility in $0.01 \mathrm{~mol} \mathrm{dm}^{-3}$ of $\mathrm{H}_2 \mathrm{SO}_4$.

pH of $0.08 \mathrm{~mol} \mathrm{~dm}^{-3} \mathrm{~HOCl}$ solution is 2.85 . Calculate its ionisation constant.

Calculate the pH of a solution formed by mixing equal volumes of two solutions A and B of a strong acid having $\mathrm{pH}=6$ and $\mathrm{pH}=4$ respectively.

The solubility product of $\mathrm{Al}(\mathrm{OH})_3$ is $2.7 \times 10^{-11}$. Calculate its solubility in $\mathrm{gL}^{-1}$ and also find out pH of this solution. (Atomic mass of $\mathrm{Al}=27 \mathrm{u}$ )

Calculate the volume of water required to dissolve 0.1 g lead (II) chloride to get a saturated solution.

$\left(\mathrm{K}_{\text {sp }}\right.$ of $\mathrm{PbCl}_2=3.2 \times 10^{-8}$, atomic mass of $\left.\mathrm{Pb}=207 \mathrm{u}\right)$

A reaction between ammonia and boron trifluoride is given below.

$$: \mathrm{NH}_3+\mathrm{BF}_3 \longrightarrow \mathrm{H}_3 \mathrm{~N}: \mathrm{BF}_3$$

Identify the acid and base in this reaction. Which theory explains it? What is the hybridisation of $B$ and $N$ in the reactants?

$$ \begin{aligned} & \text { Following data is given for the reaction } \\ & \mathrm{CaCO}_3(\mathrm{~s}) \longrightarrow \mathrm{CaO}(\mathrm{s})+\mathrm{CO}_2(\mathrm{~g}) \\ & \Delta_f H^{\ominus}[\mathrm{CaO}(s)]=-635.1 \mathrm{~kJ} \mathrm{~mol}^{-1} \\ & \Delta_f H^{\ominus}\left[\mathrm{CO}_2(g)\right]=-393.5 \mathrm{~kJ} \mathrm{~mol}^{-1} \\ & \Delta_f H^{\ominus}\left[\mathrm{CaCO}_3(s)\right]=-1206.9 \mathrm{~kJ} \mathrm{~mol}^{-1} \end{aligned}$$

Predict the effect of temperature on the equilibrium constant of the above reaction.

Match the following equilibria with the corresponding condition.

For the reaction, $\mathrm{N}_2(\mathrm{~g})+3 \mathrm{H}_2(\mathrm{~g}) \rightleftharpoons 2 \mathrm{NH}_3(\mathrm{~g})$

Equilibrium constant, $\mathrm{K}_{\mathrm{c}}=\frac{\left[\mathrm{NH}_3\right]^2}{\left[\mathrm{~N}_2\right]\left[\mathrm{H}_2\right]^3}$

Some reactions are written below in Column I and their equilibrium constants in terms of $K_c$ are written in Column II. Match the following reactions with the corresponding equilibrium constant.

Match standard free energy of the reaction with the corresponding equilibrium constant.

Match the following species with the corresponding conjugate acid.

Match the following graphical variation with their description.

	A		B
A.		1.	Variation in product concentration with time
B.		2.	Reaction at equilibrium
C.		3.	Variation in reactant concentration with time

Match the Column I with Column II.

How can you predict the following stages of a reaction by comparing the value of $K_c$ and $Q_c$ ?

(i) Net reaction proceeds in the forward direction.

(ii) Net reaction proceeds in the backward direction.

(iii) No net reaction occurs.

On the basis of Le-Chatelier principle explain how temperature and pressure can be adjusted to increase the yield of ammonia in the following reaction.

$$\mathrm{N}_2(\mathrm{~g})+3 \mathrm{H}_2(\mathrm{~g}) \rightleftharpoons 2 \mathrm{NH}_3(\mathrm{~g}) \Delta \mathrm{H}=-92.38 \mathrm{~kJ} \mathrm{~mol}^{-1}$$

What will be the effect of addition of argon to the above reaction mixture at constant volume?

A sparingly soluble salt having general formula $\mathrm{A}_x^{\mathrm{p}+} \mathrm{B}_y^{\mathrm{q}-}$ and molar solubility S is in equilibrium with its saturated solution. Derive a relationship between the solubility and solubility product for such salt.

Write a relation between $\Delta \mathrm{G}$ and Q and define the meaning of each term and answer the following.

(a) Why a reaction proceeds forward when $Q

(b) Explain the effect of increase in pressure in terms of reaction quotient Q.

For the reaction, $\mathrm{CO}(\mathrm{g})+3 \mathrm{H}_2(\mathrm{~g}) \rightleftharpoons \mathrm{CH}_4(\mathrm{~g})+\mathrm{H}_2 \mathrm{O}(\mathrm{g})$