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Lithium forms monoxide, sodium forms
peroxide, the other metals form
superoxides. The superoxide O2
? ion is
stable only in the presence of large cations
such as K, Rb, Cs.

Lithium shows exceptional
behaviour in reacting directly with nitrogen of
air to form the nitride, Li3N as well
It may be noted that although lithium has
most negative E0 value (Table 10.1), its
reaction with water is less vigorous than
that of sodium which has the least negative
E0 value among the alkali metals. This
behaviour of lithium is attributed to its
small size and very high hydration energy.
Other metals of the group react explosively
with water

They also react with proton donors such
as alcohol, gaseous ammonia and alkynes.

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

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

On combustion in excess of air, lithium forms
mainly the oxide, Li2O (plus some peroxide
Li2O2), sodium forms the peroxide, Na2O2 (and
some superoxide NaO2) whilst potassium,
rubidium and caesium form the superoxides,
MO2.

The oxides and the peroxides are colourless
when pure, but the superoxides are yellow or
orange in colour.The superoxides are also
paramagnetic.

Other halides of lithium (other than lif)are soluble
in ethanol, acetone and ethylacetate; LiCl is
soluble in pyridine also.

Lithium carbonate is not so stable to heat;
lithium being very small in size polarises a
large CO3
2? ion leading to the formation of more
stable Li2O and CO2. Its hydrogencarbonate
does not exist as a solid.

Lithium unlike other alkali metals forms
no ethynide on reaction with ethyne.

Lithium nitrate when heated gives lithium
oxide, Li2O, whereas other alkali metal
nitrates decompose to give the
corresponding nitrite

The hydration enthalpies of alkaline earth
metal ions are larger than those of alkali metal
ions. Thus, compounds of alkaline earth metals
are more extensively hydrated than those of
alkali metals, e.g., MgCl2 and CaCl2 exist as
MgCl2.6H2O and CaCl2· 6H2O while NaCl and
KCl do not form such hydrates.

Calcium,strontium and barium impart characteristic
brick red, crimson and apple green colours
respectively to the flame

Beryllium and magnesium are kinetically inert
to oxygen and water because of the formation
of an oxide film on their surface

Thermal decomposition of (NH4)2BeF4 is the
best route for the preparation of BeF2

All the
elements except beryllium combine with
hydrogen upon heating to form their hydrides,
MH2.
BeH2, however, can be prepared by the reaction
of BeCl2 with LiAlH4.
2BeCl +LiAlH4?2BeH2 + LiCl + AlCl3

The alkaline
earth metals burn in oxygen to form the
monoxide, MO which, except for BeO, have
rock-salt structure

The solubility, thermal stability and the
basic character of these hydroxides increase
with increasing atomic number from Mg(OH)2
to Ba(OH)2. The alkaline earth metal
hydroxides are, however, less basic and less
stable than alkali metal hydroxides. Beryllium
hydroxide is amphoteric in nature as it reacts
with acid and alkali both.
Be(OH)2 + 2OH? ? [Be(OH)4]2?
Beryllate ion
Be(OH)2 + 2HCl + 2H2O ? [Be(OH)4]Cl2

the corresponding hydrated halides
of Be and Mg on heating suffer hydrolysis

Carbonates of alkaline earth
metals are insoluble in waterThe solubility
of carbonates in water decreases as the atomic
number of the metal ion increases.

BeSO4, and MgSO4 are readily soluble in water;
the solubility decreases from CaSO4 to BaSO4.

All of them decompose on heating to
give the oxide like lithium nitrate.
2M (NO3)2 ?2MO+4NO +O
(M = Be, Mg, Ca, Sr, Ba)

The oxide and hydroxide of beryllium,
unlike the hydroxides of other elements in
the group, are amphoteric in nature

Beryllium hydroxide dissolves in excess of
alkali to give a beryllate ion, [Be(OH)4]2? just
as aluminium hydroxide gives aluminate
ion, [Al(OH)4]?.

Problem 10.4
Why does the solubility of alkaline earth
metal hydroxides in water increase down
the group?
Solution
Among alkaline earth metal hydroxides,
the anion being common the cationic
radius will influence the lattice enthalpy.
Since lattice enthalpy decreases much
more than the hydration enthalpy with
increasing ionic size, the solubility
increases as we go down the group.
Problem 10.5
Why does the solubility of alkaline earth
metal carbonates and sulphates in water
decrease down the group?
Solution
The size of anions being much larger
compared to cations, the lattice enthalpy
will remain almost constant within a
particular group. Since the hydration
enthalpies decrease down the group,
solubility will decrease as found for
alkaline earth metal carbonates and
sulphates.

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