NCERT Solutions

A liquid is in equilibrium with its vapour in a sealed container at a fixed temperature. The volume of the container is suddenly increased. a) What is the initial effect of the change on vapour pressure? b) How do rates of evaporation and condensation change initially? c) What happens when equilibrium is restored finally and what will be the final vapour pressure?

What is $K_{\mathrm{c}}$ for the following equilibrium when the equilibrium concentration of each substance is: $[\mathrm{SO}_{2}]=0.60\mathrm{M},[\mathrm{O}_{2}]=0.82\mathrm{M}$ and $[\mathrm{SO}_{3}]=1.90\mathrm{M}$ ? $2 \mathrm{SO}_{2}(\mathrm{~g})+\mathrm{O}_{2}(\mathrm{~g}) \rightleftharpoons 2 \mathrm{SO}_{3}(\mathrm{~g})$

At a certain temperature and total pressure of $10^{5} \mathrm{~Pa}$, iodine vapour contains $40\%$ by volume of I atoms $\mathrm{I}_{2}(\mathrm{~g}) \rightleftharpoons 2 \mathrm{I}(\mathrm{~g})$

Calculate $K_{p}$ for the equilibrium.

For the following equilibrium, $K_{c}=6.3 \times 10^{14}$ at $1000 \mathrm{~K}$ $\mathrm{NO}(\mathrm{g})+\mathrm{O}_{3}(\mathrm{~g}) \rightleftharpoons \mathrm{NO}_{2}(\mathrm{~g})+\mathrm{O}_{2}(\mathrm{~g})$ Both the forward and reverse reactions in the equilibrium are elementary bimolecular reactions. What is $K_{c}$, for the reverse reaction?

Explain why pure liquids and solids can be ignored while writing the equilibrium constant expression?

Reaction between $\mathrm{N}_{2}$ and $\mathrm{O}_{2-}$ takes place as follows: $2 \mathrm{~N}_{2}(\mathrm{~g})+\mathrm{O}_{2}(\mathrm{~g}) \rightleftharpoons 2 \mathrm{~N}_{2} \mathrm{O}(\mathrm{~g})$

If a mixture of $0.482 \mathrm{~mol} \mathrm{~N}_{2}$ and $0.933 \mathrm{~mol}$ of $\mathrm{O}_{2}$ is placed in a $10 \mathrm{~L}$ reaction vessel and allowed to form $\mathrm{N}_{2} \mathrm{O}$ at a temperature for which $K_{\mathrm{c}}=2.0 \times 10^{-37}$, determine the composition of equilibrium mixture.

Nitric oxide reacts with $\mathrm{Br}_{2}$ and gives nitrosyl bromide as per reaction given below: $2 \mathrm{NO}(\mathrm{g})+\mathrm{Br}_{2}(\mathrm{~g}) \rightleftharpoons 2 \mathrm{NOBr}(\mathrm{g})$

When $0.087 \mathrm{~mol}$ of $\mathrm{NO}$ and $0.0437 \mathrm{~mol}$ of $\mathrm{Br}_{2}$ are mixed in a closed container at constant temperature, $0.0518 \mathrm{~mol}$ of $\mathrm{NOBr}$ is obtained at equilibrium. Calculate equilibrium amount of $\mathrm{NO}$ and $\mathrm{Br}_{2}$.

At $450 \mathrm{~K}, K_{p}=2.0 \times 10^{10} / \mathrm{bar}$ for the given reaction at equilibrium. $2 \mathrm{SO}_{2}(\mathrm{~g})+\mathrm{O}_{2}(\mathrm{~g}) \rightleftharpoons 2 \mathrm{SO}_{3}(\mathrm{~g})$

What is $K_{c}$ at this temperature ?

A sample of $\mathrm{HI}(\mathrm{g})$ is placed in flask at a pressure of $0.2 \mathrm{~atm}$. At equilibrium the partial pressure of $\mathrm{HI}(\mathrm{g})$ is $0.04 \mathrm{~atm}$. What is $K_{p}$ for the given equilibrium ? $2 \mathrm{HI}(\mathrm{g}) \rightleftharpoons \mathrm{H}_{2}(\mathrm{~g})+\mathrm{I}_{2}(\mathrm{~g})$

A mixture of $1.57 \mathrm{~mol}$ of $\mathrm{N}_{2}$, $1.92 \mathrm{~mol}$ of $\mathrm{H}_{2}$ and $8.13 \mathrm{~mol}$ of $\mathrm{NH}_{3}$ is introduced into a $20 \mathrm{~L}$ reaction vessel at $500 \mathrm{~K}$. At this temperature, the equilibrium constant, $K_{\mathrm{c}}$ for the reaction $\mathrm{N}_{2}(\mathrm{~g})+3 \mathrm{H}_{2}(\mathrm{~g}) \rightleftharpoons 2 \mathrm{NH}_{3}(\mathrm{~g})$ is $1.7 \times 10^{2}$. Is the reaction mixture at equilibrium? If not, what is the direction of the net reaction?

The equilibrium constant expression for a gas reaction is, $K_{c}=\frac{\left[\mathrm{NH}_{3}\right]^{4}\left[\mathrm{O}_{2}\right]^{5}}{[\mathrm{NO}]^{4}\left[\mathrm{H}_{2} \mathrm{O}\right]^{6}}$

Write the balanced chemical equation corresponding to this expression.

One mole of $\mathrm{H}_{2} \mathrm{O}$ and one mole of $\mathrm{CO}$ are taken in $10 \mathrm{~L}$ vessel and heated to $725 \mathrm{~K}$. At equilibrium $40\%$ of water (by mass) reacts with $\mathrm{CO}$ according to the equation, $\mathrm{H}_{2} \mathrm{O}(\mathrm{g})+\mathrm{CO}(\mathrm{g}) \rightleftharpoons \mathrm{H}_{2}(\mathrm{~g})+\mathrm{CO}_{2}(\mathrm{~g})$

Calculate the equilibrium constant for the reaction.

At $700 \mathrm{~K}$, equilibrium constant for the reaction: $\mathrm{H}_{2}(\mathrm{~g})+\mathrm{I}_{2}(\mathrm{~g}) \rightleftharpoons 2 \mathrm{HI}(\mathrm{g})$ is $54.8$. If $0.5 \mathrm{~mol} \mathrm{~L}^{-1}$ of $\mathrm{HI}(\mathrm{g})$ is present at equilibrium at $700 \mathrm{~K}$, what are the concentration of $\mathrm{H}_{2}(\mathrm{~g})$ and $\mathrm{I}_{2}(\mathrm{~g})$ assuming that we initially started with $\mathrm{HI}(\mathrm{g})$ and allowed it to reach equilibrium at $700 \mathrm{~K}$ ?

What is the equilibrium concentration of each of the substances in the equilibrium when the initial concentration of ICl was $0.78 \mathrm{~M}$ ? $2 \mathrm{ICl}(\mathrm{g}) \rightleftharpoons \mathrm{I}_{2}(\mathrm{~g})+\mathrm{Cl}_{2}(\mathrm{~g}) ; \quad K_{c}=0.14$

$K_{p}=0.04 \mathrm{~atm}$ at $899 \mathrm{~K}$ for the equilibrium shown below. What is the equilibrium concentration of $\mathrm{C}_{2} \mathrm{H}_{6}$ when it is placed in a flask at $4.0 \mathrm{~atm}$ pressure and allowed to come to equilibrium? $\mathrm{C}_{2} \mathrm{H}_{6}(\mathrm{~g}) \rightleftharpoons \mathrm{C}_{2} \mathrm{H}_{4}(\mathrm{~g})+\mathrm{H}_{2}(\mathrm{~g})$

A sample of pure $\mathrm{PCl}_{5}$ was introduced into an evacuated vessel at $473 \mathrm{~K}$. After equilibrium was attained, concentration of $\mathrm{PCl}_{5}$ was found to be $0.5 \times 10^{-1} \mathrm{~mol} \mathrm{~L}^{-1}$. If value of $K_{\mathrm{c}}$ is $8.3 \times 10^{-3}$, what are the concentrations of $\mathrm{PCl}_{3}$ and $\mathrm{Cl}_{2}$ at equilibrium? $\mathrm{PCl}_{5}(\mathrm{~g}) \rightleftharpoons \mathrm{PCl}_{3}(\mathrm{~g})+\mathrm{Cl}_{2}(\mathrm{~g})$

One of the reaction that takes place in producing steel from iron ore is the reduction of iron(II) oxide by carbon monoxide to give iron metal and $\mathrm{CO}_{2}$. $\mathrm{FeO}(\mathrm{s})+\mathrm{CO}(\mathrm{g}) \rightleftharpoons \mathrm{Fe}(\mathrm{s})+\mathrm{CO}_{2}(\mathrm{~g}) ; K_{p}=0.265 \text { atm at } 1050 \mathrm{~K}$ What are the equilibrium partial pressures of $\mathrm{CO}$ and $\mathrm{CO}_{2}$ at $1050 \mathrm{~K}$ if the initial partial pressures are: $p_{\mathrm{CO}}=1.4 \mathrm{~atm}$ and $p_{\mathrm{CO}_{2}}=0.80 \mathrm{~atm}$ ?

Equilibrium constant, $K_{c}$ for the reaction $\mathrm{N}_{2}(\mathrm{~g})+3 \mathrm{H}_{2}(\mathrm{~g}) \rightleftharpoons 2 \mathrm{NH}_{3}(\mathrm{~g}) \text { at } 500 \mathrm{~K} \text { is } 0.061$

At a particular time, the analysis shows that composition of the reaction mixture is $3.0 \mathrm{~mol} \mathrm{~L}^{-1} \mathrm{~N}_{2}$, $2.0 \mathrm{~mol} \mathrm{~L}^{-1} \mathrm{H}_{2}$ and $0.5 \mathrm{~mol} \mathrm{~L}^{-1} \mathrm{NH}_{3}$. Is the reaction at equilibrium? If not in which direction does the reaction tend to proceed to reach equilibrium?

Bromine monochloride, BrCl decomposes into bromine and chlorine and reaches the equilibrium: $2 \mathrm{BrCl}(\mathrm{g}) \rightleftharpoons \mathrm{Br}_{2}(\mathrm{~g})+\mathrm{Cl}_{2}(\mathrm{~g})$ for which $K_{\mathrm{c}}=32$ at $500 \mathrm{~K}$. If initially pure BrCl is present at a concentration of $3.3 \times 10^{-3} \mathrm{~mol} \mathrm{~L}^{-1}$, what is its molar concentration in the mixture at equilibrium?

At $1127 \mathrm{~K}$ and $1 \mathrm{~atm}$ pressure, a gaseous mixture of $\mathrm{CO}$ and $\mathrm{CO}_{2}$ in equilibrium with soild carbon has $90.55\% \mathrm{CO}$ by mass $\mathrm{C}(\mathrm{s})+\mathrm{CO}_{2}(\mathrm{~g}) \rightleftharpoons 2 \mathrm{CO}(\mathrm{~g})$

Calculate $K_{\mathrm{c}}$ for this reaction at the above temperature.

Calculate a) $\Delta G^{\ominus}$ and b) the equilibrium constant for the formation of $\mathrm{NO}_{2}$ from $\mathrm{NO}$ and $\mathrm{O}_{2}$ at $298 \mathrm{~K}$ $\mathrm{NO}(\mathrm{g})+1 / 2 \mathrm{O}_{2}(\mathrm{~g}) \rightleftharpoons \mathrm{NO}_{2}(\mathrm{~g})$ where

The equilibrium constant for the following reaction is $1.6 \times 10^{5}$ at $1024 \mathrm{~K}$ $\mathrm{H}_{2}(\mathrm{~g})+\mathrm{Br}_{2}(\mathrm{~g}) \rightleftharpoons 2 \mathrm{HBr}(\mathrm{~g})$ Find the equilibrium pressure of all gases if $10.0 \mathrm{~bar}$ of $\mathrm{HBr}$ is introduced into a sealed container at $1024 \mathrm{~K}$.

Describe the effect of: a) addition of $\mathrm{H}_2$ b) addition of $\mathrm{CH}_3\mathrm{OH}$ c) removal of $\mathrm{CO}$ d) removal of $\mathrm{CH}_3\mathrm{OH}$ on the equilibrium of the reaction: $2 \mathrm{H}_{2}(\mathrm{~g})+\mathrm{CO}(\mathrm{g}) \quad \rightleftharpoons \quad \mathrm{CH}_{3} \mathrm{OH}(\mathrm{g})$

At $473 \mathrm{~K}$, equilibrium constant $K_{c}$ for decomposition of phosphorus pentachloride, $\mathrm{PCl}_{5}$ is $8.3 \times 10^{-3}$. If decomposition is depicted as, $\mathrm{PCl}_{5}(\mathrm{~g}) \quad \rightleftharpoons \mathrm{PCl}_{3}(\mathrm{~g})+\mathrm{Cl}_{2}(\mathrm{~g}) \quad \Delta_{\mathrm{r}} H^{\ominus}=124.0 \mathrm{~kJ} \mathrm{~mol}^{-1}$ a) write an expression for Kc for the reaction. b) what is the value of Kc for the reverse reaction at the same temperature? c) what would be the effect on $K_{\mathrm{c}}$ if (i) more $\mathrm{PCl}_{5}$ is added (ii) pressure is increased (iii) the temperature is increased ?

Dihydrogen gas used in Haber's process is produced by reacting methane from natural gas with high temperature steam. The first stage of two stage reaction involves the formation of $\mathrm{CO}$ and $\mathrm{H}_{2}$. In second stage, $\mathrm{CO}$ formed in first stage is reacted with more steam in water gas shift reaction, $\mathrm{CO}(\mathrm{g})+\mathrm{H}_{2} \mathrm{O}(\mathrm{g}) \rightleftharpoons \mathrm{CO}_{2}(\mathrm{~g})+\mathrm{H}_{2}(\mathrm{~g})$ If a reaction vessel at $400^{\circ} \mathrm{C}$ is charged with an equimolar mixture of $\mathrm{CO}$ and steam such that $p_{\mathrm{co}}=p_{\mathrm{H}_{2} \mathrm{O}}=4.0$ bar, what will be the partial pressure of $\mathrm{H}_{2}$ at equilibrium? $K_{p}=10.1$ at $400^{\circ} \mathrm{C}$

Predict which of the following reaction will have appreciable concentration of reactants and products: a) $\mathrm{Cl}_2(\mathrm{~g}) \rightleftharpoons 2 \mathrm{Cl}(\mathrm{g}) \quad K_c=5 \times 10^{-39}$ b) $\mathrm{Cl}_2(\mathrm{~g})+2 \mathrm{NO}(\mathrm{g}) \rightleftharpoons 2 \mathrm{NOCl}(\mathrm{g}) \quad K_c=3.7 \times 10^{8}$ c) $\mathrm{Cl}_2(\mathrm{~g})+2 \mathrm{NO}_{2}(\mathrm{~g}) \rightleftharpoons 2 \mathrm{NO}_{2} \mathrm{Cl}(\mathrm{g}) \quad K_{c}=1.8$

The value of $K_{c}$ for the reaction $3 \mathrm{O}_{2}(\mathrm{~g}) \rightleftharpoons 2 \mathrm{O}_{3}(\mathrm{~g})$ is $2.0 \times 10^{-50}$ at $25^{\circ} \mathrm{C}$. If the equilibrium concentration of $\mathrm{O}_{2}$ in air at $25^{\circ} \mathrm{C}$ is $1.6 \times 10^{-2}$, what is the concentration of $\mathrm{O}_{3}$ ?

The reaction, $\mathrm{CO}(\mathrm{g})+3 \mathrm{H}_{2}(\mathrm{~g}) \rightleftharpoons \mathrm{CH}_{4}(\mathrm{~g})+\mathrm{H}_{2} \mathrm{O}(\mathrm{~g})$ is at equilibrium at $1300 \mathrm{~K}$ in a $1 \mathrm{~L}$ flask. It also contain $0.30 \mathrm{~mol}$ of $\mathrm{CO}$, $0.10 \mathrm{~mol}$ of $\mathrm{H}_{2}$ and $0.02 \mathrm{~mol}$ of $\mathrm{H}_{2} \mathrm{O}$ and an unknown amount of $\mathrm{CH}_{4}$ in the flask. Determine the concentration of $\mathrm{CH}_{4}$ in the mixture. The equilibrium constant, $K_{c}$ for the reaction at the given temperature is $3.90$.

What is meant by the conjugate acid-base pair? Find the conjugate acid/base for the following species: $\mathrm{HNO}_{2}, \mathrm{CN}^{-}, \mathrm{HClO}_{4}, \mathrm{~F}^{-}, \mathrm{OH}^{-}, \mathrm{CO}_{3}^{2-}$, and $\mathrm{S}^{2-}$

Which of the followings are Lewis acids? $\mathrm{H}_{2} \mathrm{O}, \mathrm{BF}_{3}, \mathrm{H}^{+}$, and $\mathrm{NH}_{4}^{+}$

What will be the conjugate bases for the Brönsted acids: $\mathrm{HF}, \mathrm{H}_{2} \mathrm{SO}_{4}$ and $\mathrm{HCO}_{3}^{-}$?

Write the conjugate acids for the following Brönsted bases: $\mathrm{NH}_{2}^{-}, \mathrm{NH}_{3}$ and $\mathrm{HCOO}^{-}$.

The species: $\mathrm{H}_{2} \mathrm{O}, \mathrm{HCO}_{3}^{-}, \mathrm{HSO}_{4}^{-}$ and $\mathrm{NH}_{3}$ can act both as Brönsted acids and bases. For each case give the corresponding conjugate acid and base.

The concentration of hydrogen ion in a sample of soft drink is $3.8 \times 10^{-3} \mathrm{M}$. What is its $\mathrm{pH}$ ?

The $\mathrm{pH}$ of a sample of vinegar is $3.76$. Calculate the concentration of hydrogen ion in it.

The ionization constant of $\mathrm{HF}, \mathrm{HCOOH}$ and $\mathrm{HCN}$ at $298 \mathrm{~K}$ are $6.8 \times 10^{-4}$, $1.8 \times 10^{-4}$ and $4.8 \times 10^{-9}$ respectively. Calculate the ionization constants of the corresponding conjugate base.

The ionization constant of phenol is $1.0 \times 10^{-10}$. What is the concentration of phenolate ion in $0.05 \mathrm{~M}$ solution of phenol? What will be its degree of ionization if the solution is also $0.01 \mathrm{~M}$ in sodium phenolate?

The first ionization constant of $\mathrm{H}_{2} \mathrm{~S}$ is $9.1 \times 10^{-8}$. Calculate the concentration of $\mathrm{HS}^{-}$ ion in its $0.1 \mathrm{~M}$ solution. How will this concentration be affected if the solution is $0.1 \mathrm{~M}$ in $\mathrm{HCl}$ also? If the second dissociation constant of $\mathrm{H}_{2} \mathrm{~S}$ is $1.2 \times 10^{-13}$, calculate the concentration of $\mathrm{S}^{2-}$ under both conditions.

The ionization constant of acetic acid is $1.74 \times 10^{-5}$. Calculate the degree of dissociation of acetic acid in its $0.05 \mathrm{~M}$ solution. Calculate the concentration of acetate ion in the solution and its $\mathrm{pH}$.

It has been found that the $\mathrm{pH}$ of a $0.01 \mathrm{~M}$ solution of an organic acid is $4.15$. Calculate the concentration of the anion, the ionization constant of the acid and its $\mathrm{p} K_{\mathrm{a}}$.

Calculate the $\mathrm{pH}$ of the following solutions: a) $2 \mathrm{~g}$ of $\mathrm{TlOH}$ dissolved in water to give $2 \mathrm{~litre}$ of solution. b) $0.3 \mathrm{~g}$ of $\mathrm{Ca}(\mathrm{OH})_{2}$ dissolved in water to give $500 \mathrm{~mL}$ of solution. c) $0.3 \mathrm{~g}$ of $\mathrm{NaOH}$ dissolved in water to give $200 \mathrm{~mL}$ of solution. d) $1 \mathrm{~mL}$ of $13.6 \mathrm{~M} \mathrm{HCl}$ is diluted with water to give $1 \mathrm{~litre}$ of solution.

The degree of ionization of a $0.1 \mathrm{~M}$ bromoacetic acid solution is $0.132$. Calculate the $\mathrm{pH}$ of the solution and the $p K_{\mathrm{a}}$ of bromoacetic acid.

The $\mathrm{pH}$ of $0.005 \mathrm{~M}$ codeine $\left(\mathrm{C}_{18} \mathrm{H}_{21} \mathrm{NO}_{3}\right)$ solution is $9.95$. Calculate its ionization constant and $\mathrm{p} K_{\mathrm{b}}$.

What is the $\mathrm{pH}$ of $0.001 \mathrm{~M}$ aniline solution? The ionization constant of aniline can be taken from Table 6.7. Calculate the degree of ionization of aniline in the solution. Also calculate the ionization constant of the conjugate acid of aniline.

The ionization constant of dimethylamine is $5.4 \times 10^{-4}$. Calculate its degree of ionization in its $0.02 \mathrm{~M}$ solution. What percentage of dimethylamine is ionized if the solution is also $0.1 \mathrm{~M}$ in $\mathrm{NaOH}$ ?

The $\mathrm{pH}$ of milk, black coffee, tomato juice, lemon juice and egg white are $6.8, 5.0, 4.2, 2.2$ and $7.8$ respectively. Calculate corresponding hydrogen ion concentration in each.

If $0.561 \mathrm{~g}$ of $\mathrm{KOH}$ is dissolved in water to give $200 \mathrm{~mL}$ of solution at $298 \mathrm{~K}$. Calculate the concentrations of potassium, hydrogen and hydroxyl ions. What is its $\mathrm{pH}$ ?

The solubility of $\mathrm{Sr}(\mathrm{OH})_{2}$ at $298 \mathrm{~K}$ is $19.23 \mathrm{~g} / \mathrm{L}$ of solution. Calculate the concentrations of strontium and hydroxyl ions and the $\mathrm{pH}$ of the solution.

The ionization constant of propanoic acid is $1.32 \times 10^{-5}$. Calculate the degree of ionization of the acid in its $0.05 \mathrm{~M}$ solution and also its $\mathrm{pH}$. What will be its degree of ionization if the solution is $0.01 \mathrm{~M}$ in $\mathrm{HCl}$ also?

The $\mathrm{pH}$ of $0.1 \mathrm{~M}$ solution of cyanic acid (HCNO) is $2.34$. Calculate the ionization constant of the acid and its degree of ionization in the solution.

The ionization constant of nitrous acid is $4.5 \times 10^{-4}$. Calculate the $\mathrm{pH}$ of $0.04 \mathrm{~M}$ sodium nitrite solution and also its degree of hydrolysis.

A $0.02 \mathrm{~M}$ solution of pyridinium hydrochloride has $\mathrm{pH}=3.44$. Calculate the ionization constant of pyridine.

Predict if the solutions of the following salts are neutral, acidic or basic: $\mathrm{NaCl}, \mathrm{KBr}, \mathrm{NaCN}, \mathrm{NH}_{4} \mathrm{NO}_{3}, \mathrm{NaNO}_{2}$ and $\mathrm{KF}$

The ionization constant of chloroacetic acid is $1.35 \times 10^{-3}$. What will be the $\mathrm{pH}$ of $0.1 \mathrm{~M}$ acid and its $0.1 \mathrm{~M}$ sodium salt solution?

Ionic product of water at $310 \mathrm{~K}$ is $2.7 \times 10^{-14}$. What is the $\mathrm{pH}$ of neutral water at this temperature?

Calculate the $\mathrm{pH}$ of the resultant mixtures: a) $10 \mathrm{~mL}$ of $0.2 \mathrm{~M} \mathrm{Ca}(\mathrm{OH})_{2}+25 \mathrm{~mL}$ of $0.1 \mathrm{~M} \mathrm{HCl}$ b) $10 \mathrm{~mL}$ of $0.01 \mathrm{~M} \mathrm{H}_{2} \mathrm{SO}_{4}+10 \mathrm{~mL}$ of $0.01 \mathrm{~M} \mathrm{Ca}(\mathrm{OH})_{2}$ c) $10 \mathrm{~mL}$ of $0.1 \mathrm{~M} \mathrm{H}_{2} \mathrm{SO}_{4}+10 \mathrm{~mL}$ of $0.1 \mathrm{~M} \mathrm{KOH}$

Determine the solubilities of silver chromate, barium chromate, ferric hydroxide, lead chloride and mercurous iodide at $298 \mathrm{~K}$ from their solubility product constants given in Table 6.9. Determine also the molarities of individual ions.

The solubility product constant of $\mathrm{Ag}_{2} \mathrm{CrO}_{4}$ and $\mathrm{AgBr}$ are $1.1 \times 10^{-12}$ and $5.0 \times 10^{-13}$ respectively. Calculate the ratio of the molarities of their saturated solutions.

Equal volumes of $0.002 \mathrm{~M}$ solutions of sodium iodate and cupric chlorate are mixed together. Will it lead to precipitation of copper iodate? (For cupric iodate $K_{\mathrm{sp}}=7.4 \times 10^{-8}$ ).

The ionization constant of benzoic acid is $6.46 \times 10^{-5}$ and $K_{\mathrm{sp}}$ for silver benzoate is $2.5 \times 10^{-13}$. How many times is silver benzoate more soluble in a buffer of $\mathrm{pH} 3.19$ compared to its solubility in pure water?

What is the maximum concentration of equimolar solutions of ferrous sulphate and sodium sulphide so that when mixed in equal volumes, there is no precipitation of iron sulphide? (For iron sulphide, $K_{\mathrm{sp}}=6.3 \times 10^{-18}$ ).

What is the minimum volume of water required to dissolve $1 \mathrm{~g}$ of calcium sulphate at $298 \mathrm{~K}$ ? (For calcium sulphate, $K_{\mathrm{sp}}$ is $9.1 \times 10^{-6}$ ).

The concentration of sulphide ion in $0.1 \mathrm{~M} \mathrm{HCl}$ solution saturated with hydrogen sulphide is $1.0 \times 10^{-19} \mathrm{~M}$. If $10 \mathrm{~mL}$ of this is added to $5 \mathrm{~mL}$ of $0.04 \mathrm{~M}$ solution of the following: $\mathrm{FeSO}_{4}, \mathrm{MnCl}_{2}, \mathrm{ZnCl}_{2}$ and $\mathrm{CdCl}_{2}$. in which of these solutions precipitation will take place?