The s- block elements of the periodic table are those in which the last electron enters the outermost
s-orbital. As the s-orbital can accommodate only two electrons, two group (1 & 2) belong to the s-block of
the periodic table. Group 1 of the periodic table consists of the elements : lithium (Li), sodium (Na), potassium
(K), rubidium (Rb), caesium (Cs) and francium (Fr). They are collectively known as the alkali metals. These
are so called because they form hydroxides on reaction with water which are strongly alkaline in nature.
All the alkali metal have one valence electron, ns1 outside the noble gas core. The single valence electron
is at a long distance from the nucleus and is only weakly held. Hence the loosely held s-electrons in the
outermost valence shell of these elements make them the most electropositive 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 alkali metal atoms have the largest size in a particular period of the periodic table. The atomic and
ionic radii of alkali metal increase on moving down the group i.e., they increase in size while going from
Li to Cs.
All the alkali metal are silvery white, soft (because of having only one valency electron which participate in
bonding) and light metals. Because of the larger size, these element have low density which increases
down the group from Li to Cs.
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 superoxide.
The alkali metals react with dihydrogen at about 673 K (lithium at 1073 K) to form hydrides. All the metal
hydrides are ionic solid with high melting points
The alkali metals readily react vigorously with halogens to form ionic halide , MX. However lithium halides
are some what covalent. It is because of the high polarisation capability of lithium ion
The alkali metals, are strong reducing agents, lithium being the most and sodium the least powerful with
the small size of its ion, lithium has the highest hydration enthalpy which account for its high negative Eï‘
value and its high reducing power
In dilute solutions the main species are metal ions (M+) and electrons, which are solvated (i.e. ammoniated).
The blue colour, corresponding to a broad absorption band near 1500 nm that falls into the visible range, is
attributed to the solvated electron.
The anomalous behavior of lithium is due to the : (i) exceptionally small size of its atom and ion, and (ii) high
polarising power (i.e., charge/ radius ratio ). As a result, there is increased covalent character of lithium
compound which is responsible for their solubility in organic solvent. Further, lithium shows diagonal
relationship to magnesium
The similarity between lithium and magnesium is particularly striking and arises because of their
similar size: atomic radii, Li = 152 pm, Mg = 160 pm; ionic radii : Li+ = 76 pm, Mg2+ = 72 pm. The
main points of similarity are :
Beryllium the first member of the Group 2 metals, shows anomalous behaviour as compared to
magnesium and rest of the members. Further, it shows diagonal relationship to aluminium
The ionic radius of Be2+ is estimated to be 31 pm; the charge/radius ratio is nearly the same as that
of the Al3+ ion. Hence beryllium resembles aluminium in some ways. Some of the similarities are
By reduction of nitrites and nitrates of sodium with metallic sodium :
It is formed by heating sodium in excess of air free from moisture and carbon dioxide or in excess
of pure oxygen
It is prepared by burning potassium in excess of oxygen free from moisture.
PREPARATION : It is most conveniently manufactured by one of the following processes.
(a) Methods involving sodium carbonate as a starting material
1 It is a white crystalline solid and has soapy touch.
2 It’s density is 2.13 g/mL and melting point is 318.4°C.
It is prepared by electrolysis of KCl solution.
KOH resembles NaOH in all its reactions. However KOH is much more soluble in alcohol. This accounts
for the use of alcoholic KOH in organic chemistry.
Step - 1 (In ammonia absorber (A)) : Saturation of brine with ammonia and CO2
The potassium carbonate like sodium carbonate, can not be prepared by Solvay process because the
intermediate, KHCO3 formed is soluble in appreciable amount in water.
It is obtained as the intermediate product in the Solvay ammonia soda process. Normal carbonate can be
changed to bicarbonate by passing carbon dioxide through its saturated solution.
1 It is a white crystalline solid and effloresces readily in dry air to form anhydrous sodium sulphate.
2 It is reduced to sodium sulphide when heated with carbon
Magnesium oxide when mixed with a saturated solution of magnesium chloride sets to a hard
mass like cement known as magnesia cement or Sorel’s cement. The composition is
MgCl2.5MgO.xH2O.
It is a white powder. It is sparingly soluble in water. It is basic in nature and forms salt with acids. It
decomposes on heating. It readily dissolves in ammonium chloride solution and is, therefore, not precipitated
in group IIIrd of qualitative analysis
PROPERTIES :
It is a colourless crystalline solid, highly deliquescent and highly soluble in water
3Ca(OH)2 + 2Cl2  Ca(OCl)2. Ca(OH)2. CaCl2. 2H2O (bleaching powder).
It can be prepared by adding sodium bicarbonate to a hot solution of magnesium salt
It can be obtained by passing carbon dioxide through lime water or by adding sodium carbonate solution to
CaCl2.
It occurs in nature as minerals kiesserite (MgSO4.H2O), epsom salt (MgSO4.7H2O)and kainite
(KCl.MgSO4.3H2O).
It is found in nature as anhydride (CaSO4) and gypsum (CaSO4.2H2O)
It can be prepared by reacting any calcium salt with either sulphuric acid or a soluble sulphate
PREPARATION :
It is obtained when gypsum, calcium sulphate dihydrate (CaSO4.2H2O), is heated at 120°C (393 K).
