1. Electronic Configuration
The silvery lustre of alkali metals is due to the presence of highly mobile electrons of the metallic lattice. There being only a single electron per atom, the metallic bonding is not so strong. As a result of this, these metals are soft in nature. However, the softness increases with increases of atomic number because there is continuous decrease of metallic bond strength on account of increase in atomic size.
2. Atomic and ionic radii
At radii (Å)
All are light metals. The sensities of metals Li, Na and K Are lighter than water. Density gradually increases in moving down from Li to Cs. Potassium is however, lighter than sodium. Sequence of densities: Li < Na < K < Rb < Cs
4. Structures of the metals, hardness and cohesive energy
5. Melting and boiling points
The energy binding the atoms in the crystal lattices of these metals in relatively low on account of a single electron in the valency shell. Consequently, the metals have low melting and boiling points. These decrease in moving down from Li to Cs as the metallic bond strength decreases or cohesive force decrease.
6. Ionisation energies and electropositive character
Due to their large size, the outermost electron is far from the nucleus and can easily be removed. Their ionization potentials are relatively low. Thus, the metals have a great tendency to lose the ns1 electron to change into M + ions. These metals are highly electropositive in nature. As the ionization potential decreases from Li to Cs, the electropositive character increases, i.e. metallic character increases. The reactivity of these metals increases from Li to Cs.
Ionization Potential (eV)
7. Oxidation States
The alkali metals can lose their ns1 electron quite easily to form univalent positive ion, M+. The ion has a stable configuration of an inert gas. These metals are univalent in nature and show electro valency, i.e., form electro valent compounds.
The alkali metals are good conductors of heat and electricity. This is due to the presence of loosely held valency electrons which are free to move throughout the metal structure.
10. Heat of atomization
11. Flame colours & spectra
The alkali metals and their salts impart a characteristic colour to flame.
The colour actually arises from electronic transitions in short-lived species which are formed
Li Na K Rb Cs
Momentarily in the flame Crimson red Golden red Pale violet Violet Violet
12. Reducing nature
An element, which acts as a reducing agent, must have low ionization energy. Alkali metals act as strong reducing agents as their ionization energy values are low. Since ionization energy decreases on moving down from Li to Cs, the reducing property should increases in the same order but from E°ox values it is observed that Li is the strongest reducing agent a mongst alkali metals in solution a s E°ox value of Li is maximum.
It is due to high hydration energy of Li + in a queous solutions.
Oxidation Potential (V)
1. Action with Air
On exposure to moist air, all alkali metals except lithium tarnish quickly. The effect of atmosphere increases from Li to Cs. These are, therefore always kept under kerosene oil to protect them from air.
2. Action with Water
Alkali metals decompose water with the evolution of hydrogen gas.
3. Action with hydrogen The alkali metals combine directly with hydrogen to form crystalline hydrides of the formula M. H. These hydrides are ionic and contain the hydrides ion. H–
2M +H2 ——→ 2MH
The ionic character of the bonds in these hydride increases from LiH to CsH and the stability decreases in the same order. They are powerful reducing agents especially at high temperatures.
SiCl4 + 4NaH ——→ SiH4 + 4NaCl
4. Action with Oxygen (Oxides and Hydroxides)
The normal oxides ‘M2O’ react with water to form hydroxides.
M2O + H2O ——→ 2MOH
The basic nature of the oxides (M2O) increases gradually on moving down in the group. The hydroxides (MOH) are colourless, strong alkaline and corrosive compounds. These are soluble in water and dissolve with evolution of heat. The hydroxides are thermally stable except LiOH. The relative strength of the hydroxides increases from LiOH to CsOH.
CsOH > RbOH > KOH > NaOH > LiOH
5. Action with halogens The alkali metals directly react with halogens forming the halides of the type MX. 2M + X2 → 2MX
With the exception of certain lithium haloids, the metal halides are ionic compounds (M+X–). The halides are crystalline and have high melting and boiling points. The fused halides are good conductors of electricity and are u sed for the preparation of alkali metals. All halides except LiF dissolve in water.
KI + I2 → K[I3], KBr + ICI → K[BrICI], KF + BrF3 → K[BrF4]
6. Natural of Oxysalts
Alkali metals readily react with oxyacids forming corresponding salts with evolution of hydrogen. Lithium salt behave abnormally due to polarizing power of Li+ ion (small size) and lattice energy effects.
7. Natural of carbonates and bicarbonates
As the electropositive character increases from Li to Cs, the solubility of the carbonates increases in the same order
Cs2CO3 > Rb2CO3 > K2CO3 > Na2CO3 > Li2CO3
Li2CO3 decomposes on heating and is insoluble in water. The aqueous solution of carbonates are alkaline. This is due to hydrolysis as carbonates are salts of strong bases and weak acid (H2CO3) carbonic acid).
The bicarbonates, MHCO3, of the alkali metals, with the exception of lithium, are known in solid state. The bicarbonates are soluble in water. On heating, bicarbonates decompose into carbonates with evolution of CO2.
2MHCO3 ——→ +M2CO3 + H2O + CO2
The abnormal behaviour of Li2CO3 towards heat can be explained in the following manner. The Li + ion exerts a strong polarizing action and distorts the electron cloud of the nearby atom the large CO32- ion. This results in the weak ening of the C—O bond and strengthening of the Li—O bond. This ultimately facilitates the decomposition of Li2CO3 into Li2O and CO2. The lattice energy Li2O is higher than the lattice energy of carbonate. This also favours the decomposition of Li2CO3.
8. Natural of nitrates
Nitrates of the type, MNO3, are known. These are colourless, soluble in water and electrovalent in nature. The nitrates do not undergo hydrolysis. With the exception of LiNO3, the other nitrates decompose to nitrites and oxygen.
2MNO3 ——→ 2MNO2 + O2
Lithium nitrate decomposes to oxide on heating, it is due to diagonal relationship with magnesium.
9. Nature of Sulphates
Sulphates of the type M2SO4 are known. With the exception of Li2SO4, other sulphates are solu ble in water. The sulphates when fused with carbon form sulphides.
M2SO4 + 4C ——→ M2S + 4CO
The sulphates of alkali metals form double salts with the sulphates of the trivalent metals like Fe, Al, Cr, etc. The double sulphates crystallize with large number of water molecules are potash alum K2SO4, Al2(SO4)3.24H2O consists of 24 water molecules. Sulphate of lithium are not known to form alum.
10. Action of liquid ammonia
The alkali metals dissolve in liquid ammonia without the evolution of hydrogen. The colour of the dilute solutions is blue. The metal atom loses electron and it combines with ammonia molecule.
M → M+ (in liquid ammonia) + e (ammoniated)
M+(x + y) NH3 → [M(NH3)x]+ + e(NH3)y solvated electron
On heating its blue colour changes to bronze. It is ammoniated electron which is responsible for colour.
The solutions are good conductors of electricity and have strong reducing properties. The solutions are paramagnetic in nature.
When dry ammonia is passed over hot metal, amides are formed.
The amides are decomposed by cold water with evolution of NH3 MNH2 + 2NH3 ——→ MOH + NH3
Recent studies proved the existence of Li(NH3)4, a golden yellow solid.
11. Formation of alloys
The alkali metals form alloys amongst themselves and with other metals. These combine with mercury readily forming amalgams.
12. Complex Formation
Alkali metals have a very little tendency to form complexes, Lithium being small in size forms certain complexes but this tendency decreases as the size increases.
DIAGONAL RELATIONSHIP: SIMILARITIES WITH MAGNESIUM
Lithium shows resemblance with magnesium, an element of group IIA. This resemblance is termied as diagonal relationship.
Reason for the diagonal relationship are the following:
(i) Electronegativities of Li and Mg are quite comparable (Li = 1.00 Mg = 1.20)
(ii) Atomic ra dii and ionic ra dii of Li and Mg are not very mu ch different.
Atomic radii (Å) Li-1.23, Mg-1.36
Ionic radii (Å) Li+-00.60, Mg2+-0.65
(iii) Atomic volumes of Li and Mg are go to similar Li-12.97 mL, Mg-13.97 mL
(iv) Both have high polarizing power (ionic potential) Polarizing power
Cations with large ionic potentials have a tendency to polarize the anions and to give partial covalent character to compounds.
Lithium resembles magnesium in the following respects.
(a) Both Li and Mg are harder and have higher melting points than the other metals in their respective groups.
(b) Li like Mg decomposes water slowly to liberate hydrogen.
2Li + 2H2O ——→ 2LiOH + H2; Mg + 2H2O ——→ Mg(OH)2 + H2
(c) Both elements combine with nitrogen on heating.
6Li + N2 ——→ 2Li3N; 3Mg + N2 ——→ Mg3 + N2
Both the nitrides are decomposed by water with evolution of ammonia (NH3)
Li3N + 3H2O ——→ 3LiOH + NH3; Mg3N2 + 6H2O ——→ 3Mg(OH)2 + 2NH3
(d) Both Li and Mg combine with carbon on heating
2Li + 3H2O ——→ 3Li2C2; Mg + 2C ——→ MgC2
Both the carbides yield C2H2 with water.
(e) Lithium forms only monoxide when heated in oxygen.
Mg also forms the monoxide
4Li + O2 ——→ 2Li2O; 2Mg + O2 ——→ 2MgO
Both the are less soluble in water.
(f) Hydroxides of Li and Mg are weak bases and are slightly soluble in water. Both decomposes on heating
2LiOH ——→ Li2O + H2O; Mg(OH)2 ——→ MgO + H2O
(g) Lithium fluoride, phosphate, oxalate and carbonate lik e the corresponding salts of Mg, are sparingly soluble in water.
(h) Carbonates of Li and Mg decomposes on heating.
Li2CO3 ——→ Li2O + CO2; MgCO3 ——→ MgO + CO2
(i) Nitrates of Li and Mg decompose on heating giving mixture of nitrogen dioxide and oxygen.
4Li2 NO3 ——→ 2Li2O + 4NO2 + O2 ; 2Mg(NO3)2 ——→ 2MgO + 4NO2 + O2
COMPOUNDS OF SODIUM
1. Sodium Oxide,
Na2O 2NaNO3 + 10 Na ——→ 6Na 2O + N2 ; 2NaNO2 + 6 Na ——→ 4Na 2O + N2
NO2O + H2O ——→ 2NaOH
2. Sodium Peroxide, Na2O2
It is formed by heating sodium in excess of air free form moisture and carbon dioxide or in excess of pure oxygen.
Properties2Na 2O2 + 2H2O ——→ 4NaOH + O2; Na 2O2 + H2SO4 ——→ Na 2SO4 + H2O2
3. Sodium Hydroxide (Caustic Soda), NaOH
It is one of the important chemicals and manufactured on a very large scale forming an important chemical industry. It is most conveniently manufactured by one of the following processes:
(a) Methods involving sodium carbonate as a starting material: Two methods are used. These are
(i) Causticisation process (Gossage process)
Na2Co3 + Ca(Oh)2 ⇆ CaCo3 + 2NaOH
(ii) Lowig’s process
Na2CO3 + Fe2O3 ——→ 2NaFe2 + CO2; 2Na2FeO2 + H2O ——→ 2NaOH + Fe2O3
(b) Methods involving sodium chloride as starting material (electrolysis of brine)
At Cathod 2H2O + 2e ⇆ H2 + 2OH- ; Na+ + OH- ⇆ NaOH
1. Strong alkali: NaOH ⇆ Na+ + OH-
(i) NaOH + HCl ——→ NaCl + H2O;
(ii) 2NaOH + CO2 ——→ Na2CO3 + H2O
(iii) Al2O3 + 2NaOH ——→ 2NaAlO2 + H2O
2. Action on non-metals Cl2 + 2NaOH ——→ NaCl + NaClO + H2O
3Cl2 + 6NaOH ——→ 5NaCl + NaClO3 + 3H2O
P4 + 3NaOH + 3H2 O ——→ 3NaH2PO2 + PH3
3. Action on Metals Zn + 2NaOH ——→ Na2ZnO2 + H2;
4. Action on Salts
(i) Ni(NO3)2 + 2NaOH ——→ Ni(OH)2 + 2NaNO3
(ii) Insoluble hydroxides which dissolve in excess of NaOH
ZnSO4 + 2NaOH ——→ Zn(OH)2 + Na2SO4
(iii) Unstable hydroxides 2AgNO3 + 2NaOH ——→ 2AgOH + 2NaNO3
2AgOH ——→ Ag2O + H2O
SODIUM CARBONATE OR WASHING SODA (Na2CO3.10H2O)
(a) Le-Blanc Process
(i) Conversion of NaCl into
Na2CO4 NaCl + H2SO4 ——→ NaHSO4 + HCl
NaHSO4 + NaCl ——→ Na2SO4 + HCl
(b) Solvay ammonia soda process
NH3 + H2O + CO2 ——→ NH4HCO3
NaCl + NH4HCO3 ——→ NaHCO3 + NH4Cl