Why does water show high boiling point as compared to hydrogen sulphide? Give reasons for your answer.
Water show high boiling point as compared to hydrogen sulphide due to high electronegativity of oxygen ( $\mathrm{EN}=3.5$ ). Water undergoes extensive H - bonding as a result of which water exists as associated molecule.
For breaking these hydrogen bond, a large amount of energy is needed and hence the boiling point of $\mathrm{H}_2 \mathrm{O}$ is high.
In other words, due to lower electronegativity of $S(E N=2.5)$, hydrogen sulphide do not undergo H -bonding. Consequently, $\mathrm{H}_2 \mathrm{~S}$ exists as discrete molecule and hence its boiling point is much lower than that of $\mathrm{H}_2 \mathrm{O}$. Thats why $\mathrm{H}_2 \mathrm{~S}$ is a gas at room temperature.
Why can dilute solutions of hydrogen peroxide not be concentrated by heating? How can a concentrated solution of hydrogen peroxide be obtained?
Dilute solutions of $\mathrm{H}_2 \mathrm{O}_2$ cannot be concentrated by heating because it decomposes much below its boiling point.
$$2 \mathrm{H}_2 \mathrm{O}_2 \longrightarrow 2 \mathrm{H}_2 \mathrm{O}+\mathrm{O}_2$$
$1 \% \mathrm{H}_2 \mathrm{O}_2$ is extracted with water and concentrated to $\sim 30 \%$ (by mass) by distillation under reduced pressure. It can be further concentrated to $\sim 85 \%$ by careful distillation under low pressure. The remaining water can be frozen out to obtain pure $\mathrm{H}_2 \mathrm{O}_2$.
Why is hydrogen peroxide stored in wax lined bottles?
Hydrogen peroxide is decomposed by rough surfaces of glass, alkali oxides present in it and light to form $\mathrm{H}_2 \mathrm{O}$ and $\mathrm{O}_2$.
$$2 \mathrm{H}_2 \mathrm{O}_2 \longrightarrow 2 \mathrm{H}_2 \mathrm{O}+\mathrm{O}_2$$
To prevent this decomposition, hydrogen peroxide is usually stored in paraffin wax coated plastic or teflon bottles.
Why does hard water not form lather with soap?
Hard water contains salts of calcium and magnesium ions. Hard water does not give lather with soap and forms scum/precipitate with soap. Soap containing sodium stearate $\left(\mathrm{C}_{17} \mathrm{H}_{35} \mathrm{COONa}\right)$ reacts with hard water to precipitate out as $\mathrm{Ca} / \mathrm{Mg}$ stearate.
$$2 \mathrm{C}_{17} \mathrm{H}_{35} \mathrm{COONa}(a q)+\mathrm{M}^{2+}(\mathrm{aq}) \longrightarrow\left(\mathrm{C}_{17} \mathrm{H}_{35} \mathrm{COO}\right)_2 \mathrm{M} \downarrow+2 \mathrm{Na}^{+}(a q)$$ (where, $M$ is $\mathrm{Ca} / \mathrm{Mg}$ )
It is therefore, unsuitable for laundry.
Phosphoric acid is preferred over sulphuric acid in preparing hydrogen peroxide from peroxides. Why?
$\mathrm{H}_2 \mathrm{SO}_4$ acts as a catalyst for decomposition of $\mathrm{H}_2 \mathrm{O}_2$. Therefore, some weaker acids such as $\mathrm{H}_3 \mathrm{PO}_4, \mathrm{H}_2 \mathrm{CO}_3$ is preferred over $\mathrm{H}_2 \mathrm{SO}_4$ for preparing $\mathrm{H}_2 \mathrm{O}_2$ from peroxides.
$$3 \mathrm{BaO}_2+2 \mathrm{H}_3 \mathrm{PO}_4 \longrightarrow \underset{\text { (insoluble) }}{\mathrm{Ba}_3\left(\mathrm{PO}_4\right)_2}+3 \mathrm{H}_2 \mathrm{O}_2 $$