Group 1 Elements (Alkali Metals)

Group 1 (Alkali Metals):

The alkali metals show regular trends in their physical and chemical properties with the increasing atomic number. Alkali metals includes the following elements:

Electronic Configuration:

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:

The alkali metal atoms have the largest sizes in a particular period of the periodic table. The atom becomes larger with increase in atomic number. 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:

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.

Hydration Enthalpy:

· The hydration enthalpies of alkali metal ions decrease with increase in ionic sizes.

Li⁺ > Na⁺ > K⁺ > Rb⁺ > Cs⁺ Li⁺

· They has maximum degree of hydration and for this reason lithium salts are mostly hydrated.e.g., LiCl. 2H₂O.

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.

· 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.

· Alkali metals can 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:

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 ion is stable only in the presence of large cations such as K, Rb, Cs.

4 Li + O₂ + Li₂O (oxide)

2. Reactivity towards Water:

· The alkali metals react with water to form hydroxide and dihydrogen.

2M + 2H₂O 2M⁺ + 2OH^- + H₂ (M = an alkali metal)

· Its reaction with water is less vigorous than that of sodium.

· This behaviour of lithium is attributed to its small size and very high hydration energy.

· Other metals of the group react explosively with water.

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^–.

· Lithium halides are somewhat covalent because of the high polarisation capability of lithium ion.

· 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.

5. Reducing nature:

· The alkali metals are strong reducing agents, lithium being the most and sodium the least powerful.

M (s)		M (g)	Sublimation enthalpy
M (g)		M^+ (g) + e^-	Ionization enthalpy
M^+ (g) + H2O		M^+ (aq)	Hydration enthalpy

6. Solutions in liquid ammonia:

· 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.

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

Uses:

· 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 or Pb alloy.

· 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.