The alkali metals show regular trends in their physical and chemical properties with the increasing atomic number. The atomic, physical and chemical properties of alkali metals are discussed below.
10.1.1 Electronic Configuration
All the alkali metals have one valence electron, ns1 (Table 10.1) outside the noble gas core. The loosely held s-electron in the outermost valence shell of these elements makes them the most electropositive metals. They readily lose electron to give monovalent M+ ions. Hence they are never found in free state in nature.
10.1.2 Atomic and Ionic Radii
The alkali metal atoms have the largest sizes in a particular period of the periodic table. With increase in atomic number, the atom becomes larger.
The monovalent ions (M+) are smaller than the parent atom. The atomic and ionic radii of alkali metals increase on moving down the group i.e., they increase in size while going from Li to Cs.
10.1.3 Ionization Enthalpy
The ionization enthalpies of the alkali metals are considerably low and decrease down the group from Li to Cs. This is because the effect of increasing size outweighs the increasing nuclear charge, and the outermost electron is very well screened from the nuclear charge.
10.1.4 Hydration Enthalpy
The hydration enthalpies of alkali metal ions decrease with increase in ionic sizes.
Li+> Na+ > K+ > Rb+ > Cs+
Li+ has maximum degree of hydration and for this reason lithium salts are mostly hydrated, e.g., LiCl· 2H2O
10.1.5 Physical Properties
All the alkali metals are silvery white, soft and light metals. Because of the large size, these elements have low density which increases down the group from Li to Cs.
However, potassium is lighter than sodium. The melting and boiling points of the alkali metals are low indicating weak metallic bonding due to the presence of only a single valence electron in them.
The alkali metals and their salts impart characteristic colour to an oxidizing flame. This is because the heat from the flame excites the outermost orbital electron to a higher energy level.
When the excited electron comes back to the ground state, there is emission of radiation in the visible region of the spectrum.
Alkali metals can therefore, be detected by the respective flame tests and can be determined by flame photometry or atomic absorption spectroscopy.
These elements when irradiated with light, the light energy absorbed may be sufficient to make an atom lose electron. This property makes caesium and potassium
useful as electrodes in photoelectric cells.
10.1.6 Chemical Properties
The alkali metals are highly reactive due to their large size and low ionization enthalpy. The reactivity of these metals increases down the group.
1-Reactivity towards air: The alkali metals tarnish in dry air due to the formation of their oxides which in turn react with moisture to form hydroxides. They burn vigorously in oxygen forming oxides. Lithium forms monoxide, sodium forms peroxide, the other metals form superoxides. The superoxide O2– ion is stable only in the presence of large cations such as K, Rb, Cs.
4Li + O2 → 2Li2O (Oxide)
2Na + O2 →Na2O2 (peroxide)
M + O2 → MO2 (superoxide) (M= K, Rb, Cs)
In all these oxides the oxidation state of the alkali metal is +1. Lithium shows exceptional behaviour in reacting directly with nitrogen of air to form the nitride, Li3N as well.
Because of their high reactivity towards air and water, alkali metals are normally kept in kerosene oil.
2-Reactivity towards water: The alkali metals react with water to form hydroxide and dihydrogen.
2 M + 2H2O → 2M+ + 2OH- + H2 (M = an alkali metal)
It may be noted that although lithium has most negative E- value, its reaction with water is less vigorous than that of sodium which has the least negative E- value among the alkali metals.
This behaviour of lithium is attributed to its small size and very high hydration energy. Other metals of the group react explosively with water. They also react with proton donors such as alcohol, gaseous ammonia and alkynes.
3-Reactivity towards dihydrogen: The alkali metals react with dihydrogen at about 673K (lithium at 1073K) to form hydrides. All the alkali metal hydrides are ionic solids with high melting points.
2M +H2 → 2M+H-
4-Reactivity towards halogens : The alkali metals readily react vigorously with halogens to form ionic halides, M+X–. However, lithium halides are somewhat covalent.
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 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.
Reducing nature: The alkali metals are strong reducing agents, lithium being the most and sodium the least powerful. The standard electrode potential (E-) which measures the reducing power represents the overall change:
M(s) → M(g) sublimationenthalpy
M(g) → M+ (g) + e- onizationenthalpy
M+(g) + H2O → M+ (aq) hydrationenthalpy
With the small size of its ion, lithium has the highest hydration enthalpy which accounts for its high negative E- value and its high reducing power.
5-Solutions in liquid ammonia: The alkali metals dissolve in liquid ammonia giving deep blue solutions which are conducting in nature.
M + (x+y) NH3 → [M(NH3)x]+ + [e(NH3)y]-
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+ (am) + e- + NH3 (l) → MNH2(am) + 1/2 H2(g)
(where ‘am’ denotes solution in ammonia.)
In concentrated solution, the blue colour changes to bronze colour and becomes diamagnetic.
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. Lithium is also used to make electrochemical cells. Sodium is used to make a Na/Pb alloy needed to make PbEt4 and PbMe4.
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. Potassium has a vital role in biological systems.
Potassium chloride is used as a fertilizer. Potassium hydroxide is used in the manufacture of soft soap. It is also used as an excellent absorbent of carbon dioxide. Caesium is used in devising photoelectric cells.