Group 1 Elements Alkali Metals



Group-I elements have one electron in their valence shell. They do not occur in the native or free state. These elements are collectively known as alkali metals because their oxides and hydroxides form strong alkalies like NaOH, KOH, etc. Lithium is known as bridge element.


Electronic Configuration of alkali metals

All the alkali metals have one valence electron, ns1 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.

Atomic and Ionic Radii of alkali metals

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.

Ionization Enthalpy of alkali metals

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.


The second ionisation enthalpies of all the alkali metals are very high because by releasing an electron, ions acquire noble gas configuration. so removal of second electron is difficult.

Hydration Enthalpy of alkali metals

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


Physical Properties of alkali metals

 All the alkali metals are silvery white, soft and light metals.


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 as given below:

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.


Chemical Properties of alkali metals

 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 of alkali metals  towards air

 On exposure to moist air, their surface get tarnished due to the formation of their oxides. hydroxides and carbonates.


Hence. they are kept under inert liquid like kerosene oil but lithium is kept wrapped in paraffin wax because it floats on the surface of kerosene oil due to its low density.

Note Fire due to alkali Metals is extinguished by CCI4


Action of oxygen

(a) All the alkali metals when heated with oxygen form different types of oxides. e.g., lithium forms lithium oxide (Li2O), sodium forms sodium peroxide (Na2O2), while K, Rb and Cs form superoxides MO2 (where, M = K, Rb or Cs)

The stability of peroxides and superoxides increases as the size of alkali metal increases.

(b) Superoxides are coloured and paramagnetic as these possess three electron bondwhere one unpaired electron is present.

(c) All oxides. peroxides and superoxides are basic in nature.

Basic strength of oxides increase in the order

Li2O < Na2O < K2O < Cs2O

Na2O2 acquires yellow colour due to the presence of superoxides as an impurity.

KO2 (potassium superoxide) is used as a source of oxygen in submarines, space shuttles and in emergency breathing apparatus such as oxygen masks.

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.

The reactivity of alkali metals towards hydrogen is

Li > Na > K > Rb > Cs.


Reactivity of alkali metals  towards water

2M + 2H2O → 2MOH + H2 (where, M = Li, Na, K, Rb, and Cs)

The reactivity order with water is

Li < Na < K < Rb < Cs

This is due to increase in electropositive character in the same order.

KOH is stronger base than NaOH .

LiOH is used to remove carbon dioxide from exhaled air in confined quarters like submarines and space vehicles.


Reactivity of alkali metals  towards dihydrogen


Reactivity of alkali metals  towards halogens

Alkali metals combine readily with halogens to form ionic halides

[where M= Li,Na, K etc. and X = F,Cl, Br,I]

The reactivity of alkali metals towards a particular halogen increases in the order  :

Li < Na  < K < Rb < Cs

while that of halogen  towards a particular alkali metal decreases in the order :


All alkali halides except LiF are freely soluble in water (LiF is soluble in non-polar solvents. Since it has strong covalent bond.)

The power of the cation to polarise the anion  is known as the polarising power while the tendency of the anion to get polarised is known as its polarisability. The polarising power of cation and polarisability of anion depends on the following factors (which are collectively referred to as Fajan’s rules)

Size of the cation - Smaller the size of cation greater is its polarising power. So LiCl is more covalent than KCl.

Size of the anion - Bigger the anion, larger is its polarisability. Hence the covalent character of lithium halides  is in the order -

LiI > LiBr > LiCl > LiF

Charge of the ion and electronic configuration - Larger the charge on the cation, greater is its polarising power

Thus the covalent character of various halides is in the order

when two cations have same charge and size, the one having 18 electrons in their outermost shell will have larger polarising power than a cation having 8 electrons in the outermost shell. For example CuCl is more covalent than NaCl.

Above rules help to predict the ionic /covalent character of metal halides.


Reducing nature of alkali metals

The alkali metals are strong reducing agents, lithium being the most and sodium the least powerful. The standard electrode potential (E0) which measures the reducing power represents the overall change :


With the small size of its ion, lithium has the highest hydration enthalpy which accounts for its high negative E0 value and its high reducing power.


Solutions in liquid ammonia of alkali metals

The alkali metals dissolve in liquid ammonia giving deep blue solutions which are conducting in nature.


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.


(where ‘am’ denotes solution in ammonia.)

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


Uses of alkali metals

  • 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.