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Chapter: 11th Chemistry : Alkali and Alkaline Earth Metals

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Alkali metals

The word “alkali” is derived from the word al-qalīy meaning the plant ashes, referring to the original source of alkaline substances.

Alkali metals:

The word “alkali” is derived from the word al-qalīy meaning the plant ashes, referring to the original source of alkaline substances. A water-extract of burnt plant ashes, called potash contain mainly potassium carbonate. Alkali metal group consists of the elements: lithium, sodium, potassium, rubidium, caesium and francium. They are all metals, generally soft and highly reactive. They form oxides and hydroxides and these compounds are basic in nature.


General characteristics of alkali metals:


Alkali metals are highly reactive and are found in nature only as compounds. Rubidium and caesium are found associated in minute quantities with minerals of other alkali metals. Francium is radioactive and does not occur appreciably in nature. Francium is highly radioactive; its longest-lived isotope has a half-life of only 21 minutes.


Electronic configuration

The general valence shell electronic configuration of alkali metals is ns1, where ‘n’ represents the period number.

Common oxidation state

All these elements are highly electropositive in nature. They readily lose their valence electron to give monovalent cations (M+). Alkali metals have only one oxidation state which is +1.

Atomic and ionic radii

Being the first element of each period, alkali metals have the largest atomic and ionic radii in their respective periods. On moving down the group, there is an increase in the number of shells and, therefore, atomic and ionic radii increase. The monovalent ions (M+) are smaller than the respective parent atoms as expected.

Ionisation enthalpy

Alkali metals have the lowest ionisation enthalpy compared to other elements present in the respective period. As we go down the group, the ionisation enthalpy decreases due to the increase in atomic size. In addition, the number of inner shells also increases, which in turn increases the magnitude of screening effect and consequently, the ionisation enthalpy decreases down the group.


The second ionisation enthalpies of alkali metals are very high. The removal of an electron from the alkali metals gives monovalent cations having stable electronic configurations similar to the noble gas. Therefore, it becomes very difficult to remove the second electron from the stable configurations already attained.

Hydration enthalpy

Lithium salts are more soluble than the salts of other metals of group 1. eg. LiClO4 is up to 12 times more soluble than NaClO4. KClO4, RbClO4 and CsClO4 have solubilities only 10-3 times of that of LiClO4 . The high solubility of Li salts is due to strong solvation of small size of Li+ ion.


Alkali metals have comparatively smaller value of electronegativity than the other elements in the respective period. When they react with other elements, they usually produce ionic compounds. For example, they react with halogens to form ionic halides.

Flame colour and the spectra:

When the alkali metal salts moistened with concentrated hydrochloric acid are heated on a platinum wire in a flame, they show characteristic coloured flame as shown below.

The heat in the flame excites the valence electron to a higher energy level. When it drops back to its actual energy level, the excess energy is emitted as light, whose wavelength is in the visible region as shown in the above table.


Distinctive behavior of lithium


The distinctive behaviour of Li+ ion is due to its exceptionally small size, high polarising power, high hydration energy and non availability of d-orbitals.

Diagonal Relationship:

Similarity between the first member of group 1 (Li) and the diagonally placed second element of group 2 (Mg) is called diagonal relationship. It is due to similar size (r Li+ = 0.766 Å and Mg2+ = 0.72 Å) and comparable electronegativity values (Li = 1.0; Mg = 1.2).


Chemical properties of alkali metals


Alkali metals exhibit high chemical reactivity. The reactivity of alkali metals increases from Li to Cs, since the ionisation energy decreases down the group. All alkali metals are highly reactive towards the more electronegative elements such as oxygen and halogens. Some characteristic chemical properties of alkali metals are described blow.


Reaction with oxygen


All the alkali metals on exposure to air or oxygen burn vigorously, forming oxides on their surface. Lithium forms only monoxide, sodium forms the monoxide and peroxide and the other elements form monoxide, peroxide, and superoxides. These oxides are basic in nature.


4Li +O  2Li2O (simple oxide)


2Na +O2 Na2O2  (peroxide)


M + O2  → MO2

(M= K, Rb,Cs; MO2 -superoxide)


Reaction with hydrogen


All alkali metals react with hydrogen at about 673 K (lithium at 1073 K) to form the corresponding ionic hydrides. Reactivity of alkali metals with hydrogen decreases from Li to Cs.


2M + H2 2 M+H-

(M = Li, Na, K, Rb, Cs)


The ionic character of the hydrides increases from Li to Cs and their stability decreases. The hydrides behave as strong reducing agents and their reducing nature increases down the group.


Reaction with halogen

Alkali metals combine readily with halogens to form ionic halides MX. Reactivity of alkali metals with halogens increases down the group because of corresponding decrease in ionisation enthalpy.


2M + X→ 2 MX

(M= Li, Na, K, Rb, Cs) (X= F, Cl, Br, I)


All metal halides are ionic crystals. However Lithium iodide shows covalent character, as it is the smallest cation that exerts high polarising power on the iodide anion. Additionally, the iodide ion being the largest can be polarised to a greater extent by Li+ ion.


Reaction with liquid ammonia:

Alkali metals dissolve in liquid ammonia to give deep blue solutions that are conducting in nature. The conductivity is similar to that of pure metals (The specific conductivity of Hg is 104 Ω-1 and for sodium in liquid ammonia is 0.5 x 104 Ω-1). This happens because the alkali metal atom readily loses its valence electron in ammonia solution. Both the cation and the electron are ammoniated to give ammoniated cation and ammoniated electron.


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 an amide.


M+ + e + NH3 ‚ MNH2+ ½H2


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


Reaction with water:


Alkali metals react with water to give corresponding hydroxides with the liberation of hydrogen.


2 Li + 2 H2O  → 2 LiOH+ H2


They also react with alcohol, and alkynes which contain active hydrogens.


Na + 2 C2H5OH  → 2 C2H5ONa + H2

Reducing activity:

Alkali metals can lose their valence electron readily hence they act as good reducing agents.

M(s)  → M+(g) + e

Reaction with carbon:

Lithium directly reacts with carbon to form the ionic compound, lithium carbide. Other metals do not react with carbon directly. However, when they are treated with compounds like acetylene they form acetelydes.

2 Li + 2C   → Li2C2


Uses of alkali metals:


i. Lithium metal is used to make useful alloys. For example with lead it is used 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.


ii. Lithium is also used to make electrochemical cells.


iii. Lithium carbonate is used in medicines


iv. Sodium is used to make Na/Pb alloy needed to make Pb(Et)4 and Pb(Me)4. These organolead compounds were earlier used as anti-knock additives to petrol, but nowadays lead-free petrol in use.


v. Liquid sodium metal is used as a coolant in fast breeder nuclear reactors. Potassium has vital role in biological systems.


vi. 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.


vii. Caesium is used in devising photoelectric cells.


Tags : 11th Chemistry : Alkali and Alkaline Earth Metals
Study Material, Lecturing Notes, Assignment, Reference, Wiki description explanation, brief detail