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?
Explanation for the given statement on the basis of ionisation and effect upon the concentration of sodium chloride is given below
(i) Sugar being a non-electrolyte does not ionise in water whereas NaCl ionises completely in water and produces $\mathrm{Na}^{+}$and $\mathrm{Cl}^{-}$ion which help in the conduction of electricity.
(ii) When concentration of NaCl is increased, more $\mathrm{Na}^{+}$ and $\mathrm{Cl}^{-}$ ions will be produced. Hence, conductance or conductivity of the solution increases.
$\mathrm{BF}_3$ 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?
$\mathrm{BF}_3$ is an electron deficient compound and hence acts as Lewis acid. $\ddot{\mathrm{N}} \mathrm{H}_3$ has one lone pair which it can donate to $\mathrm{BF}_3$ and form a coordinate bond. Hence, $\mathrm{NH}_3$ acts as a Lewis base.
$\mathrm{H}_3 \mathrm{N}: \longrightarrow \mathrm{BF}_3$
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?
Given that, ionisation constant of a weak base MOH
$$K_b=\left[M^{+}\right]\left[\mathrm{OH}^{-}\right][\mathrm{MOH}] .$$
Larger the ionisation constant $\left(K_b\right)$ of a base, greater is its ionisation and stronger the base. Hence, dimethyl amine is the strongest base.
$K_b\underset{5.4 \times 10^{-4}}{\text { Dimethyl amine }}>\underset{1.77 \times 10^{-5}}{\text { ammonia }}>\underset{1.77 \times 10^{-9}}{\text { pyridine }}>\underset{1.3 \times 10^{-14}}{\text { urea }}$