Chemistry Chemical properties of alkali metals

Topics to be covered

`=>` Chemical properties of alkali metals

Chemical Properties

`color{green}(★)` The alkali metals are highly reactive due to their large size and low ionization enthalpy. The reactivity of these metals increases down the group.

Reactivity towards air

`color{green}(★)` The alkali metals tarnish in dry air due to the formation of their oxides which in turn react with moisture to form hydroxides.
`color{green}(★)` They burn vigorously in oxygen forming oxides.

`color{green}(★)` Lithium forms monoxide, sodium forms peroxide, the other metals form superoxides.

`color{green}(★)` The superoxide `color{red}(O_2^(–))` ion is stable only inhe presence of large cations such as `color{red}(K, Rb, Cs)`.

`color{red}(4Li + O_2 → 2Li_2O ["oxide"])`

`color{red}(2Na +O_2 → Na_2O_2 ["peroxide"])`

`color{red}(M+O_2 → MO_2 [ "superoxide"] \ \ \ \ \ \ ( M = K , Rb, Cs))`

`color{green}(★)` In all these oxides the oxidation state of the alkali metal is +1.

`color{green}(★)` Lithium shows exceptional behaviour in reacting directly with nitrogen of air to form the nitride, `color{red}(Li_3N)` as well.

`color{green}(•)` Because of their high reactivity towards air and water, alkali metals are normally kept in kerosene oil.
Q 3282023837

What is the oxidation state of `K` in `KO_2 ?`
Class 11 Chapter ex.10 1

The superoxide species is represented as `O_2^(-)` since the compound is neutral, therefore, the oxidation state of potassium is +1.

Reactivity towards water

`color{green}(★)` The alkali metals react with water to form hydroxide and dihydrogen

`color{red}(2M + 2H_2O → 2M^(+) + 2OH^(-) + H_2 \ \ \ \ \ \ ("M = an alkali metal"))`

`color{green}(★)` It may be noted that although lithium has most negative `color{red}(E^⊖)` value (Table 10.1), its reaction with water is less vigorous than that of sodium which has the least negative `color{red}(E^⊖)` value among the alkali metals.

`color{green}(•)` This behaviour of lithium is attributed to its small size and very high hydration energy. Other metals of the group react explosively with water.

`color{green}(★)` They also react with proton donors such as alcohol, gaseous ammonia and alkynes.

Reactivity towards dihydrogen

`color{green}(★)` The alkali metals react with dihydrogen at about 673K (lithium at 1073K ) to form hydrides.

`color{green}(★)` All the alkali metal hydrides are ionic solids with high melting points.

`color{red}(2M +H_2 → 2M^+ H^(-))`

Reactivity towards halogens

`color{green}(★)` The alkali metals readily react vigorously with halogens to form ionic halides, `color{red}(M^+ X^(-))` .

`color{green}(★)` However, lithium halides are somewhat covalent.

`color{green}(•)` It is because of the high polarisation capability of lithium ion (The distortion of electron cloud of the anion by the cation is called polarisation). The `color{red}(Li^+)` ion is very small in size and has high tendency to distort electron cloud around the negative halide ion. Since anion with large size can be easily distorted, among halides, lithium iodide is the most covalent in nature.

Q 3212123930

The `E^⊖` for `Cl_2 // Cl^(-)` is +1.36, for `I_2//I^(-)` is + 0.53, for `Ag^(+) // Ag` is +0.79, `Na^(+) // Na` is `–2.71` and for `Li^(+) // Li` is – 3.04. Arrange the following ionic species in decreasing order of reducing strength : `I^(-) , Ag , Cl^(-) , Na`
Class 11 Chapter ex.10 2

The order is `Li > Na > I^(–) > Ag > Cl^(–)`

Reducing nature

`color{green}(★)` The alkali metals are strong reducing agents, lithium being the most and sodium the least powerful (Table 10.1).

`color{green}(★)` The standard electrode potential (`color{red}(E^⊖)` ) which measures the reducing power represents the overall change :

`color{red}(M(s) → M(g) \ \ \ \ \ \ \ \ \ \ \ \ ("sublimation enthalpy"))`

`color{red}(M(g) → M^(+) +e^(-) \ \ \ \ \ \ \ \ \ (" ionization enthalpy"))`

`color{red}(M^(+) (g) + H_2O → M^(+) (aq) \ \ \ \ \ \ \ (" hydration enthalpy")) `

`color{green}(★)` With the small size of its ion, lithium has the highest hydration enthalpy which accounts for its high negative `color{red}(E^⊖)` value and its high reducing power.

Solutions in liquid ammonia

`color{green}(★)` The alkali metals dissolve in liquid ammonia giving deep blue solutions which are conducting in nature.

`color{red}(M + (x+y) NH_3 → [M (NH_3)_x]^(+) + [e(NH_3)_y]^(-))`
`color{green}(★)` 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.

`color{red}(M_(am)^(+) + e^(-) + NH_3 (l) → MNH_(2(am)) + 1/2 H_2 (g))` (where ‘am’ denotes solution in ammonia.)

`color{green}(★)` In concentrated solution, the blue colour changes to bronze colour and becomes diamagnetic.


`color{green}(★)` Lithium metal is used to make useful alloys, for example with lead to make ‘white metal’ bearings for motor engines, with aluminium to make aircraft parts, and with magnesium to make armour plates. It is used in thermonuclear reactions.

`color{green}(★)` Lithium is also used to make electrochemical cells. Sodium is used to make a `color{red}(Na//Pb)` alloy needed to make `color{red}(PbEt_4)` and `color{red}(PbMe_4)`.

`color{green}(★)` These organolead compounds were earlier used as anti-knock additives to petrol, but nowadays vehicles use lead-free petrol. Liquid sodium metal is used as a coolant in fast breeder nuclear reactors.

`color{green}(★)` Potassium has a vital role in biological systems.

`color{green}(•)` Potassium chloride is used as a fertilizer.

`color{green}(•)` Potassium hydroxide is used in the manufacture of soft soap.

`color{green}(•)` It is also used as an excellent absorbent of carbon dioxide.

`color{green}(★)` Caesium is used in devising photoelectric cells.