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Inorganic Chemistry
Chapter - Gp I Alkali Metals
Lithium shows different properties due to its extremely small size and lesser metallic
character so as it shows significance similarity to Mg and other group II , which is a general
trend and can be assumed to be a general perception.
Due to high reactivity they do not occur in free state in nature the different sources
Lithium occurs a silicates [K2O.Al2O3.6SiO2] where as sodium is found as salt in sea water
NaNO3, Borax [Na2B4O7. 10H20] and Mirabilite [Na2SO4] etc. The potassium is in form of
sylvine KCl, carnalities [KCl.MgCl2.6H2O] and Felspar etc.
Rb and Cs are at low abundance, while Fr is rare as radioactive.
The PHYSICAL PROPERTIES of the group are as under
Softness increases down the group in this group, the reason can be given by a logic the kernel size
get bigger and bigger down the group and so the softness increases, this also gives a vibe that the
density increases the trend is important because of the single exception which is
Li < K < Na < Rb < Cs < Fr (density)
Melting point decreases down the group this is also because the metallic bonding is getting weaker
and weaker as we move down the group. Now the next is the Hydration enthalpy, since charge is
same on all ions the inverse relation of size is the deciding factor so,
Hydration enthalpy decreases down the group. Now some properties which are decided by the
enthalpy itself is hydrated ionic radius follows,
Li+ > K+ > Na+ > Rb+ > Cs+ > Fr+ (hydrated ionic radius) now since the hydrated size decreases down
the group the mobility is more of a smaller ion so mobility increases down the group.
Oxidation potential most of the redox reactions take place in aqueous medium and gp I metal are
good reducing agents, while Li is strongest reducing agent in aq. Medium while we don’t have a
very accurate order are trend is irregular. Keeping this point in mind will help in resolving reactions
with lithium compounds.
Flame Test .
Crimson red
golden yellow
pale violet
red violet
sky blue
The flame test is mostly asked either directly or to identify the element present in compound, while
in case of heating the Na the electron do a transaction of 3s  2s and releases photon whose
wavelength lies in the visible range causing the metal to turn pink.
Hydroxides - all hydroxides are ionic in nature. Since they are basic being metal oxides we can
conclude that more the metallic character more is basic so basicity increases down the group. Same
is logic being ionic, more metallic more ionic so solubility increases down the group.
The nature is colourless crystalline solid and now considering the thermal stability all except LiOH
are stable to heat and stability increases down the group.
LiOH on heating decomposes to  Li2O +H2O the reason here is that the oxide have a more stable
lattice than that of hydroxide of lithium
Reaction with ammonia –
M + NH3  M+(NH3)X + ammoniated electrons
The solution of ammonia due to the presence of ammoniated electrons appears to be blue. If left for
long the colour fades because of the following reaction
M+ + NH3 + e-  H2 + MNH2 the electrons get consumed and colour fades.
The solution when diluted is paramagnetic and is a good conductor of electricity for obvious reason.
But when the conc. Is high the characteristics of the solutions change The solution turns copper bronze due to the formation the metallic lustre.
 The opposite spin electron attract and solution turns diamagnetic.
 In presence of a catalyst like Fe we faster the reaction which forms MNH2.
 While all other metals are forming MNH2, here also Li shows a unique behaviour and reacts
as a group II metal to form Li2NH.
Reaction with Air –
The different metals form different oxides with oxygen
Li + O2  Li2O
(colourless oxide)
Na + O2  Na2O2 (colourless peroxide) sometimes the formation of superoxide gives a colour.
M + O2  MO2
(coloured superoxide) where M= K, Rb, Cs
The colour of the superoxide’s is due to the paramagnetic behaviour, the O2- is having two covalent
bonds and a single electron, which when move from one to other atom releases photon of visible
range giving the compounds colour, and also the paramagnetic behaviour
 The stability of peroxides and super oxides increases moving down the group.
 Li here also shows an anomalous behaviour, when react with air it is the only metal to react
with N2 present.
Li + Air  Li2O + Li3N the here also driving force is high lattice energy of product.
Li3N + H2O  LiOH + NH3 the production of ammonia makes this an important reaction.
These are reaction which are given by the second group metals and so are important.
 The Uses are important of the potassium superoxide- it is used as a absorber of the carbon
dioxide and releasing di oxygen gas, which makes it helpful for use in space capsules,
submarines to function. Also is used to know the presence the carbon dioxide both
qualitatively and quantitatively.
Reaction with Water –
Reaction is simple to give the hydroxides of respective metals
M + H2O  MOH + H2
 The reaction is highly exothermic in nature
 The reactivity is measure of metallic nature so more metallic nature and so the reactivity
increases going down the group.
Li is gentle, Na melts on the surface and on exposure to sufficient water and air catches fire,
And potassium onwards catches fire and explodes.
Due to smallest size, only Li form complexes.
Reaction with Sulphur and Phosphorus–
M + S  M2S
and all metals reacts with sulphur and may forms polysulphide in
presence of excess sulphur like M2Sn where n= 2, 3,4,5,6.
M + P  M3P
the compound are according to the valences of the atom and there for
easy to remember while the metal phosphide reacts with water to give phosgene and hydroxide.
Reaction with CarbonOnly Li reacts with carbon in alkali metals to give Li2C2 Lithium Carbide
None of the other metal directly reacts but can form carbide when heated with ethylene.
Heating Effect –
 Bicarbonates – All bicarbonates gives the same reactions on heating to decompose, into
their respective carbonates. General reaction,
heat M2CO3 + H2O + CO2
Now here we must remember that there is no difference in the product But LiHCO3 is the
only which do not exist in solid form and is aq. Form with carbon dioxide atmosphere to
prevent the reaction.
 Carbonates – Now the carbonates are product on heating Bicarbonates this states that they
are thermodynamically more stable. In Fact, all carbonates are stable to heat, except Li CO3.
heat no reaction
heat Li2O + CO2 the reaction is analogous to the decomposition of group II
carbonates decomposition, i.e. all the carbonates of Gp II decompose to give oxide and none
of them are stable to heat.
Stability order -- Li2CO3 < Na2CO3 < K2CO3 < Rb2CO3 < Cs2CO3
The difference is in polarizing power of the atoms of Gp II and Li behaves as because of size.
Li+ O- C=O the small size of Li atom increases its polarizing power as Gp II
Li+ Ohaving extra charge, more polarizing power { Fajan’s Rule }.
 Nitrates – With nitrates as well Li shows a different behaviour, all of them decompose but Li
decompose to oxidesMNO3
heat MNO2 + O2 where M= Na, K, Rb, Cs
heat Li2O + NO2 + O2
As we expect, the same type of decomposition is shown by Gp II Nitrates. google docs link
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