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Alkali metals exhibit very high chemical reactivity because—

i) Low IE1, so they form M+ easily.

ii) Low heat of atomisation so their vapour are formed easily.

iii) High heat of hydration, so a lot of energy is available to break existing bonds.

If other factors are not present then Ionisation enthalpy may be taken as criteria to measure reducing power.

Lower is the I.E more is the tendency to lose $latex \displaystyle {{e}^{-}}$ and so more is the reducing power (or Reactivity).

In Gaseous state order of ReactivityLi < Na < K < Rb < Cs

But when other factors as also present (as in aqueous medium) the reactivity is explained in terms of Standard Electrode Potentials.

\begin{array}{l} Li_{\left( {aq} \right)}^ + + {e^ - } \to Li & E_{RP}^o = - 3.05\,V\\ Na_{\left( {aq} \right)}^ + + {e^ - } \to Na & E_{^{RP}}^o = - 2.71V\\ K_{\left( {aq} \right)}^ + + {e^ - } \to k & & E_{RP}^o = - 2.925V\\ Rb_{\left( {aq} \right)}^ + + {e^ - } \to Rb & E_{RP}^o = - 2.93V\\ Cs_{\left( {aq} \right)}^ + + {e^ - } \to Cs & E_{RP}^o = - 2.927V \end{array}

Thus in aqueous medium order of reactivity of alkali metal is—

Na < K < Cs < Rb < Li


SEP is a measure of tendency of an element to lose in aqueous Medium. More negative is the , higher is the ability of element to lose and hence stronger is the reducing character.

1. Reactivity Towards Air: Li forms monoxide, Na forms peroxide, other metals form superoxides. The superoxide ion is stable only in the presence of large cation such as K, Rb, Cs.

$latex \displaystyle \begin{array}{l}4\text{Li+}{{\text{O}}_{2}}\to 2\text{L}{{\text{i}}_{\text{2}}}\text{O}\left( \text{oxide} \right)\,\,\,\,\,\,\,\text{M+}{{\text{O}}_{\text{2}}}\to \text{M}{{\text{O}}_{\text{2}}}\left( \text{Superoxide} \right)\\\text{2Na+}{{\text{O}}_{\text{2}}}\to \text{N}{{\text{a}}_{\text{2}}}{{\text{O}}_{2}}\left( \text{Peroxide} \right)\,\,\,\,\,\,\,\left( \text{M=K,}\,\text{Rb,}\,\text{Cs} \right)\end{array}$

2. Reactivity Towards Water: The alkali metals react with water to form hydroxide and dihydrogen.

$latex \displaystyle 2M+2{{H}_{2}}O\to 2MOH+{{H}_{2}}+Heat$

Although Li has most negative $latex \displaystyle {{E}^{o}}$ value but its reaction with H2O is less vigorous than that of Na which has the least $latex \displaystyle {{E}^{o}}$ value. This behaviour of Li is due to its very small size and very high hydration energy.

3. Reaction with hydrogen: All the alkali metals react with hydrogen at temperature 673 K (1073 K for Li) to form hydrides.

Li > Na > K > Rb > Cs Reactivity

Reactivity order explained on the basis of L.E. which is decreasing from LiH to CsH.

All the alkali metal hydrides are ionic solids with high M.P.

$latex \displaystyle 2M+{{H}_{2}}\xrightarrow{\Delta }2MH$

4. Reactivity Towards Halogen: All alkali metals react with halogens to form ionic halides $latex \displaystyle \left( {{M}^{+}}{{X}^{-}} \right)$

$latex \displaystyle 2M+{{X}_{2}}\to 2MX$

Order of reactivity of Alkali metals—

Li < Na < K < Rb < Cs

Order of reactivity of halogens—

F2 > Cl2 > Br > I2

5. Reducing Nature: The alkali metals are strong reducing agent, Li being the most and Na is the least powerful reducing agent.

The standard electrode potential which measures the reducing power represents the overall changes—

$latex \displaystyle {{M}_{\left( S \right)}}\to {{M}_{\left( g \right)}};\,\,\Delta {{H}_{S}}$ =Enthalpy of sublimation

$latex \displaystyle {{M}_{\left( g \right)}}\to M_{\left( g \right)}^{+}+{{e}^{-}};\,\,\,I{{E}_{1}}$ = Ionisation enthalpy

$latex \displaystyle M_{\left( g \right)}^{+}\xrightarrow{+{{H}_{2}}O}M_{\left( aq \right)}^{+};\,\,\,\,\,\,\,\Delta {{H}_{h}}$ = Enthalpy of hydration.

For all alkali metals $latex \displaystyle \Delta {{H}_{S}}$ are almost similar. IE1 for Li is highest, but this value is compensated by highly exothermic enthalpy of hydration, which is maximum for smallest Li+ ion. This highly exothermic step (III) is responsible for most negative $latex \displaystyle {{E}^{0}}$ value for Li and so Li strongest R.A. in aq. medium.

Due to small size of Li+ ion, lithium has the highest hydration enthalpy which accounts for its high negative value of $latex \displaystyle {{E}^{0}}$

$latex \displaystyle \begin{array}{l}4M+Si{{O}_{2}}\to Si+2{{M}_{2}}O\\6M+{{B}_{2}}{{O}_{3}}\to 2B+3{{M}_{2}}O\\4M+3C{{O}_{2}}\to C+2{{M}_{2}}C{{O}_{3}}\\6M+A{{l}_{2}}{{O}_{3}}\to 2Al+3{{M}_{2}}O\\2M+BeC{{l}_{2}}\to Be+2MCl\end{array}$

6. Solubility in liquid NH3: All alkali metals dissolve in liq. NH3 giving deep blue solutions which are conducting in nature.

$latex \displaystyle M\rightleftharpoons {{M}^{+}}\left( in\,liquid\,N{{H}_{3}} \right)+{{e}^{-}}$

$latex \displaystyle {{e}^{-}}+xN{{H}_{3}}\to e\left( N{{H}_{3}} \right)_{x}^{-}\to $ Ammonia solvated electron.

$latex \displaystyle \begin{array}{l}M+\left( x+y \right)N{{H}_{3}}\to {{\left[ M{{\left( N{{H}_{3}} \right)}_{x}} \right]}^{+}}+{{\left[ e{{\left( N{{H}_{3}} \right)}_{y}} \right]}^{-}}\\\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,Ammoniated\,\,\,\,\,\,\,Ammoniated\\\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,cation\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,{{e}^{-}}\end{array}$

The blue colour of the solution is due to ammoniated $latex \displaystyle {{e}^{-}}$ which absorbs energy in the visible region of light and thus imparts blue colour to the solution.

The solutions are paramagnetic.

In the presence of impurities or catalyst such as Fe (or other transition metals) the alkali metal react with NH3 to form metal amide and hydrogen.

$latex \displaystyle \begin{array}{l}{{M}^{+}}+{{e}^{-}}+N{{H}_{3}}\to MN{{H}_{2}}+\frac{1}{2}{{H}_{2}}\left( g \right)\\\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\left( Ammoniated\,solution \right)\end{array}$

In concentrated solution, the blue colour changes to bronze colour and becomes diamagnetic.