Chemistry > s Block Elements > 2.0 Alkali Metals
s Block Elements
1.0 S-Block Elements
2.0 Alkali Metals
3.0 Anamolous Behaviour of Lithium
4.0 Diagonal Relationship – similarities with magnesium
5.0 Compounds of Sodium
6.0 Alkaline Earth Metals
7.0 Diagonal Relationship – Similarities with Aluminium:
8.0 Anomalous Behaviour of Beryllium
9.0 Compounds of Calcium
2.2 Chemical properties
1. Reaction with water: The metal react with water, liberating hydrogen and forming hydroxides. $$2Na + 2{H_2}O \to 2NaOH + {H_2} \uparrow $$ Reaction becomes more vigorous as we move down the group. Their reaction with water is so exothermic that they can cause a fire or even an explosion.
2. Reaction with air: All alkali metals except lithium tarnish rapidly in dry air. Na, K, Rb and Cs form oxides whereas Li, when heated in air forms a mixture of oxides and nitrides.
3. Reaction with hydrogen: All alkali metals react with dry hydrogen at about $673\ K$ to form colourless crystalline hydrides. All the alkali metal hydrides are ionic solids with high melting point.
$$2M + {H_2}\mathop \to \limits^\Delta 2MH\left( {M = Li,Na,K,Rb\ and\ Cs} \right)$$
4. Reaction with ammonia: All alkali metals are soluble in liquid ammonia giving a solution which is paramagnetic, highly conducting and highly reducing. Dilute alkali metal ammonia solutions are blue in colour.
$${M^ + } + \left( {x + y} \right)N{H_3} \to {\left[ {M{{\left( {N{H_3}} \right)}_x}} \right]^ + } + {\left[ {e{{(N{H_3})}_y}} \right]^ - }$$
- ${\left[ {M{{\left( {N{H_3}} \right)}_x}} \right]^ + }$: Ammoniated ion cause conductivity.
- ${\left[ {e{{(N{H_3})}_y}} \right]^ - }$: ammoniated electron responsible for blue colour and paramagnetism.
The blue colour of the solution is due to the ammoniated electron which absorbs energy in the visible region of light and thus imparts blue colour to the solution. The solutions are paramagnetic and on standing slowly liberate hydrogen resulting in the formation of amide.
$$M_{\left( {am} \right)}^ + + {e^ - } + N{H_3}\left( l \right) \to MN{H_2}_{\left( {am} \right)} + \frac{1}{2}{H_2}\left( g \right)$$
(where ‘am’ denotes solution in ammonia.)
In concentrated solution, the blue colour changes to bronze colour and becomes diamagnetic.
5. Reaction with mercury: Alkali metals combine with mercury to form amalgams. The reaction is highly exothermic in nature. $$Na + Hg \to Na\left[ {Hg} \right]\left( {{\text{Sodium Amalgam}}} \right)$$
6. Reaction with sulphur and phosphorus: All metals form sulphides and phosphides.
$$\begin{equation} \begin{aligned} 3Na + P \to N{a_3}P\left( {{\text{Sodium Phosphide}}} \right) \\ 2Na + S \to N{a_2}S\left( {{\text{Sodium Sulphide}}} \right) \\\end{aligned} \end{equation} $$
7. Oxides, peroxide and superoxides:
- Li when burn in $O_2$ forms mainly lithium monoxide $(O^-2)$ $$4Li + {O_2} \to 2L{i_2}O\left( {ionic} \right)$$
- Na forms peroxide ${\left( {{O_2}} \right)^{2 - }}$ on burning in $O_2$.
- Alkali metals reacts with $O_2$ to form superoxides of the type $M{O_2}{\left( {{O_2}} \right)^ - }$. $$K + {O_2} \to K{O_2}\left( {{\text{Potassium Superoxide}}} \right)$$ All alkali metals, their oxides, peroxides and superoxides readily dissolve in water to produce corresponding hydroxides which are strong alkalies. For example: $$\begin{equation} \begin{aligned} 2Na + 2{H_2}O \to 2NaOH + {H_2} \\ N{a_2}O + 2{H_2}O \to 2NaOH + {H_2} \\ N{a_2}{O_2} + 2{H_2}O \to 2NaOH + {H_2}{O_2} \\ 2K{O_2} + 2{H_2}O \to 2KOH + {H_2}{O_2} + {O_2} \\\end{aligned} \end{equation} $$ Thus peroxides and superoxides also act as oxidizing agents since they react with $H_2O$ forming $H_2O_2$ and $O_2$ respectively. The hydroxides of all alkali metals are white crystalline solids. Two factors are responsible for hydroxide ion formation:
(i) Polarity of bond
(ii) Internuclear distance between the oxygen of the hydroxide and metal ion.
