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Alkali Metals and Their Compounds
The group 1 of the periodic table contains six
elements, namely lithium ( Li ), sodium ( Na ), potassium
( K ), rubidium ( Rb ), caesium ( Cs ) and francium ( Fr ). All
these elements are typical metals. Francium is
radioactive with longest lived isotope Fr
223
with half
life period of only 21 minute. These are usually referred
to as alkali metals since their hydroxides form strong
bases or alkalies.
(1) Electronic configuration
Elements Discovery Electronic
configuration (
1
ns )
Li
3
Arfwedson
2 1
[He ] 2 s
Na
11
Davy (1807)
10 1
[Ne] 3 s
K
19
Davy (1807)
18 1
[Ar] 4 s
Rb
37
Bunsen (1861)
36 1
[Kr] 5 s
Cs
55
Bunsen (1860)
54 1
[Xe] 6 s
Fr
87
Percy (1939)
86 1
[Rn] 7 s
(2) Occurrence : Alkali metals are very reactive
and thus found in combined state some important ores
of alkali metals are given ahead.
(i) Lithium : Triphylite, Petalite, lepidolite,
Spodumene [ LiAl ( SiO 3 ) 3 ], Amblygonite [ Li ( Al F ) PO 4 ]
(ii) Sodium : Chile salt petre ( NaNO 3
), Sodium
chloride ( NaCl ), Sodium sulphate ( Na 2 SO 4 ), Borax
( Na 2
B
4
O
7
10 H
2
O ), Glauber salt ( Na 2
SO
4
.10 H
2
O )
(iii) Potassium : Sylime ( KCl ), carnallite
( KCl. MgCl 2 .6 H 2 O ) and Felspar ( K 2 O.Al 2 O 3 .6SiO 2 )
(iv) Rubidium : Lithium ores Lepidolite, triphylite
contains 0.7 to 3% Rb 2 O
(v) Caesium : Lepidolite, Pollucite contains 0.2 to
7% Cs 2
O
(3) Extraction of alkali metals : Alkali metals
cannot be extracted by the usual methods for the
extraction of metals due to following reasons.
(i) Alkali metals are strong reducing agents,
hence cannot be extracted by reduction of their oxides
or other compounds.
(ii) Being highly electropositive in nature, it is not
possible to apply the method of displacing them from
their salt solutions by any other element.
(iii) The aqueous solutions of their salts cannot be
used for extraction by electrolytic method because
hydrogen ion is discharged at cathode instead of an
alkali metal ions as the discharge potentials of alkali
metals are high. However, by using Hg as cathode,
alkali metal can be deposited. The alkali metal readily
combines with Hg to form an amalgam from which its
recovery difficult. The only successful method,
therefore, is the electrolysis of their fused salts, usually
chlorides. Generally, another metal chloride is added to
lower their fusion temperature.
Fused NaCl :
NaCl Na Cl
fusion
offusedsaltCathode Na e Na
Electrolysis Anode Cl Cl e
: : 2 2 2
: : 2 2
2
(4) Alloys Formation
(i) The alkali metals form alloys among
themselves as well as with other metals.
s and p-Block Elements
Chapter
(ii) Alkali metals also get dissolved in mercury to
form amalgam with evolution of heat and the
amalgamation is highly exothermic.
Physical properties
(1) Physical state
(i) All are silvery white, soft and light solids.
These can be cut with the help of knife. When freshly
cut, they have bright lustre which quickly tarnishes due
to surface oxidation.
(ii) These form diamagnetic colourless ions since
these ions do not have unpaired electrons, ( i.e. M
has
ns
0
configuration). That is why alkali metal salts are
colourless and diamagnetic.
(2) Atomic and ionic radii
(i) The alkali metals have largest atomic and ionic
radii than their successive elements of other groups
belonging to same period.
(ii) The atomic and ionic radii of alkali metals,
however, increases down the group due to progressive
addition of new energy shells.
No doubt the nuclear charge also increases on
moving down the group but the influence of addition of
energy shell predominates
Li Na K Rb Cs Fr
Atomic radius (pm) 152 186 227 248 265
375
Ionic radius of M
60 95 133 148 169
ions (pm)
(3) Density
(i) All are light metals, Li , Na and K have density
less than water. Low values of density are because
these metals have high atomic volume due to larger
atomic size. On moving down the group the atomic size
as well as atomic mass both increase but increase in
atomic mass predominates over increase in atomic size
or atomic volume and therefore the ratio mass/volume
i.e. density gradually increases down the groups
(ii) The density increases gradually from Li to Cs ,
Li is lightest known metal among all.
Li = 0.534, Na = 0.972, K = 0.86, Rb = 1.53 and
Cs = 1.87 g/ml at 20
0
C.
(iii) K is lighter than Na because of its unusually
large atomic size.
(iv) In solid state, they have body centred cubic
lattice.
(4) Melting point and Boiling point
(i) All these elements possess low melting point
and boiling point in comparison to other group
members.
Li Na K Rb Cs
Fr
melting point (K) 453.5 370.8 336.2 312.
301.5 –
boiling point (K) 1620 1154.4 1038.5 961.
978.0 –
(ii) The lattice energy of these atoms in metallic
crystal lattice relatively low due to larger atomic size
and thus possess low melting point and boiling point on
moving down the group, the atomic size increases and
binding energy of their atoms in crystal lattice
decreases which results lowering of melting point.
(iii) Lattice energy decreases from Li to Cs and
thus melting point and boiling also decreases from Li to
Cs.
(5) Ionisation energy and electropositive or
metallic character
(i) Due to unpaired lone electron in ns sub-shell
as well as due to their larger size, the outermost
electron is far from the nucleus, the removal of
electron is easier and these low values of ionisation
energy. ( I.E. )
(ii) Ionisation energy of these metal decreases
from Li to Cs.
Ionisation energy Li Na K Rb Cs
Fr
IE 1
520 495 418 403 376 –
IE 2
7296 4563 3069 2650 2420 –
A jump in 2nd ionisation energy (huge
difference) can be explained as,
2 1
Re
1
2
Re
2
2 1
Li : 1 s 2 s Li : 1 s Li : 1 s
moval of
selectron
movalof
selectron
Removal of 1s electrons from Li
and that too from
completely filled configuration requires much more
energy and a jump in 2nd ionisation is noticed.
(iii) Lower are ionisation energy values, greater
is the tendency to lose ns
1
electron to change in M
ion
( i.e. M M
-
) and therefore stronger is
electropositive character.
(iv) Electropositive character increases from Li to
Cs.
Due to their strong electropositive character, they
emit electrons even when exposed to light showing
photoelectric effect. This property is responsible for
the use of Cs and K in photoelectric cell.
(6) Oxidation number and valency
(i) Alkali metals are univalent in nature due to
low ionisation energy values and form ionic
compounds. Lithium salts are, however, covalent.
(ii) Further, the M
ion acquires the stable noble
gas configuration. It requires very high values of
energy to pull out another electron from next to outer
shell of M
ion and that is why their second ionisation
energy is very high. Consequently, under ordinary
Chemical properties
(1) Formation of oxides and hydroxides
(i) These are most reactive metals and have
strong affinity for O 2
quickly tranish in air due to the
formation of a film of their oxides on the surface. These
are, therefore, kept under kerosene or paraffin oil to
protect them from air,
Peroxide
2 2
Oxide
2 2
M O MO MO
(ii) When burnt air ( O 2
), lithium forms lithium
oxide ( Li 2
O ) sodium forms sodium peroxide ( Na 2
O
2
and other alkali metals form super oxide ( Mo 2 i.e.
KO
2
, RbO 2
or CsO 2
Lithuimoxide
2 2
2 Li O LiO ;
2 2 2
2 Na O Na O
Potassium superoxide
2 2
K O KO
The reactivity of alkali metals towards oxygen to
form different oxides is due to strong positive field
around each alkali metal cation. Li
being smallest,
possesses strong positive field and thus combines with
small anion O
2 –
to form stable Li 2 O compound. The Na
and K
being relatively larger thus exert less strong
positive field around them and thus reacts with larger
oxygen anion i.e,
1
2
2
2
O and O to form stable oxides.
The monoxide, peroxides and superoxides have O 2
and
1
2
2
2
O , O ions respectively. The structures of each
are,
The O 2
ion has a three electron covalent bond
and has one electron unpaired. It is therefore
superoxides are paramagnetic and coloured KO 2
is light
yellow and paramagnetic substance.
(iii) The oxides of alkali metals and metal itself
give strongly alkaline solution in water with evolution
of heat
M HO MOH H ; H ve
2 2
Li O HO 2 LiOH ; H ve
2 2
Na O HO NaOH HO H ve
l
2 2 2 2 2 ()
KO HO KOH HO O H ve
l g
2 2 2 2 () 2 ()
The peroxides and superoxides act as strong
oxidising agents due to formation of H 2
O
2
(iv) The reactivity of alkali metals towards air
and water increases from Li to Cs that is why lithium
decomposes H 2
O very slowly at 25
o
C whereas Na does
so vigorously, K reacts producing a flame and Rb, Cs do
so explosively.
2 2
M HO MOH H
(v) The basic character of oxides and hydroxides
of alkali metals increases from Li to Cs. This is due to
the increase in ionic character of alkali metal
hydroxides down the group which leads to complete
dissociation and leads to increase in concentration of
OH
ions.
(2) Hydrides
(i) These metals combine with H to give white
crystalline ionic hydrides of the general of the formula
MH.
(ii) The tendency to form their hydrides, basic
character and stability decreases from Li to Cs since the
electropositive character decreases from Cs to Li.
2 M + H
2
2 MH ; Reactivity towards H 2
is Cs < Rb
< K < Na < Li.
(iii) The metal hydrides react with water to give
MOH and H 2 ; MH + H 2 O MOH + H 2
(iv) The ionic nature of hydrides increases from Li
to Cs because of the fact that hydrogen is present in the
these hydrides as H
and the smaller cation will
produce more polarisation of anion (according to
Fajans rule) and will develop more covalent character.
(v) The electrolysis of fused hydrides give H 2
at
anode. NaH Contains Na and Hi. e .,
fused
At cathode: Na
-
Na ; At anode:
H H e
2
2
1
(vi) Alkali metals also form hydrides like NaBH 4 ,
LiAlH 4 which are good reducing agent.
(3) Carbonates and Bicarbonates
(i) The carbonates ( M 2 CO 3 ) & bicarbonates
( MHCO
3
) are highly stable to heat, where M stands for
alkali metals.
(ii) The stability of these salts increases with the
increasing electropositive character from Li to Cs. It is
therefore Li 2 CO 3 decompose on heating, Li 2 CO 3
Li 2 O + CO 2
(iii) Bicarbonates are decomposed at relatively
low temperature,
2 3 2 2
300
3
0
2 MHCO MCO HO CO
C
(iv) Both carbonates and bicarbonates are soluble
in water to give alkaline solution due to hydrolysis of
carbonate ions or bicarbonate ions.
(4) Halides
(i) Alkali metals combine directly with halogens
to form ionic halide
M X.
(ii) The ease with which the alkali metals form
halides increases from Li to Cs due to increasing
electropositive character from Li to Cs.
(iii) Lithium halides however have more covalent
nature. Smaller is the cation, more is deformation of
[: O
O :]
1 –
Superoxide
( O 2
)
[ O O
]
2 –
x
Peroxide ( O 2
2 –
)
x
: O O
Monoxide ( O 2 )
anion and thus more is covalent nature in compound.
Also among lithium halides, lithium iodide has
maximum covalent nature because of larger anion
which is easily deformed by a cation. Thus covalent
character in lithium halides is, LiI > LiBr > LiCl > LiF
(iv) These are readily soluble in water. However,
lithium fluoride is sparingly soluble. The low solubility
of LiF is due to higher forces of attractions among
smaller Li
and smaller F
ions (high lattice energy).
(v) Halides having ionic nature have high m.pt.
and good conductor of current. The melting points of
halides shows the order, NaF > NaCl > NaBr > Nal
(vi) Halides of potassium, rubidium and caesium
have a property of combining with extra halogen atoms
forming polyhalides.
KI + I 2 KI 3 ; In KI 3(aq) the ions K
and I
3 are
present
(5) Solubility in liquid NH 3
(i) These metals dissolve in liquid NH 3
to produce
blue coloured solution, which conducts electricity to an
appreciable degree.
(ii) With increasing concentration of ammonia,
blue colour starts changing to that of metallic copper
after which dissolution of alkali metals in NH 3 ceases.
(iii) The metal atom is converted into
ammoniated metal in i.e. M
( NH 3 ) and the electron set
free combines with NH 3 molecule to produce ammonia
solvated electron.
Ammoniatedelectron
3
Ammoniatedcation
3 3
( ) [ ( )] [( )]
x y
Na x yNH NaNH eNH
(iv) It is the ammoniated electron which is
responsible for blue colour, paramagnetic nature and
reducing power of alkali metals in ammonia solution.
However, the increased conductance nature of these
metals in ammonia is due to presence of ammoniated
cation and ammonia solvated electron.
(v) The stability of metal-ammonia solution
decreases from Li to Cs.
(vi) The blue solution on standing or on heating
slowly liberates hydrogen, 2 M + 2 NH 3
2 MNH
2
+ H
2
Sodamide ( NaNH 2 ) is a waxy solid, used in preparation
of number of sodium compounds.
(6) Nitrates : Nitrates of alkali metals ( MNO 3 ) are
soluble in water and decompose on heating. LiNO 3
decomposes to give NO 2
and O 2
and rest all give
nitrites and oxygen.
2 MNO
3
2 MNO
2
+ O
2
(except Li )
4 LiNO 3 2 Li 2 O + 4 NO 2 + O 2
(7) Sulphates
(i) Alkali metals’ sulphate have the formula M 2 SO 4
(ii) Except Li 2
SO
4
, rest all are soluble in water.
(iii) These sulphates on fusing with carbon form
sulphides, M 2
SO
4
+ 4 C M
2
S + 4 CO
(iv) The sulphates of alkali metals (except Li )
form double salts with the sulphate of the trivalent
metals like Fe , Al, Cr etc. The double sulphates crystallize
with large number of water molecules as alum. e.g. K 2 SO 4.
Al 2
( SO
4
3
. 24 H
2
O.
(8) Reaction with non-metals
(i) These have high affinity for non-metals.
Except carbon and nitrogen, they directly react with
hydrogen, halogens, sulphur, phosphorus etc. to form
corresponding compounds on heating.
2 Na + H 2
C
0
300
2 NaH ; 2 K + H 2 2 KH
2 Na + Cl 2 2 NaCl ; 2 K + Cl 2 2 KCl
2 Na + S Na 2
S ; 2 K + S K
2
S
3 Na + P Na 3 P ; 3 K + P K 3 P
(ii) Li reacts, however directly with carbon and
nitrogen to form carbides and nitrides.
2 Li + 2 C LiC 2
; 6 Li + 2 N 2
2 Li 3
N
(iii) The nitrides of these metals on reaction with
water give NH 3.
M
3
N + 3 H
2
O 3 MOH + NH
3
(9) Reaction with acidic hydrogen : Alkali metals
react with acids and other compounds containing acidic
hydrogen ( i.e, H atom attached on F,O, N and triply
bonded carbon atom, for example, HF, H 2 O , ROH ,
RNH 2 , CH CH ) to liberate H 2.
2 2
2
1
M HO MOH H ;
2
2
1
M HX MX H
2
2
1
M ROH ROH H ;
2 2
2
1
M RNH RNHNa H
(10) Complex ion formation : A metal shows
complex formation only when it possesses the
following characteristics, (i) Small size (ii) High
nuclear charge (iii) Presence of empty orbitals in order
to accept electron pair ligand. Only Lithium in alkali
metals due to small size forms a few complex ions Rest
all alkali metals do not possess the tendency to form
complex ion.
Anomalous behaviour of Lithium
Anomalous behaviour of lithium is due to
extremely small size of lithium its cation on account of
small size and high nuclear charge, lithium exerts the
greatest polarizing effect out of all alkali metals on
negative ion. Consequently lithium ion possesses
remarkable tendency towards solvation and develops
covalent character in its compounds. Li differs from
other alkali metals in the following respects,
(1) It is comparatively harder than other alkali
metals. Li can’nt be stored in kerosene as it floats to
the surface, due to its very low density. Li is generally
kept wrapped in parrafin wax.
Atomic radii 1.
Ionic radii 0.60( Li
0.65( Mg
Atomic volume 12.97 c.c
13.97 c.c
(10) Both have high polarizing power.
Polarizing Power = Ionic charge / (ionic radius)
2
(11) Li and Mg Form only monooxide on heating in
oxygen.
4 Li + O 2 2 Li 2 O ; 2 Mg + O 2 2 MgO
(12) Li 2 SO 4 like MgSO 4 does not form alums.
(13) The bicarbonates of Li and Mg do not exist in
solid state, they exist in solution only.
(14) Alkyls of Li and Mg (R. Li and R. MgX ) are
soluble in organic solvent.
(15) Lithium chloride and MgCl 2 both are
deliquescent and separate out from their aqueous
solutions as hydrated crystals, LiCl. 2 H 2
O and MgCl 2
2 H 2 O.
Uses of Lithium
(1) It is used as a deoxidiser in metallurgy of Cu
and Ni.
(2) It is used as an alloying metal with
(i) Pb to give toughened bearings.
(ii) Al to give high strength Al - alloy for aircraft
industry.
(iii) Mg (14% Li ) to give extremely tough and
corrosion resistant alloy which is used for armour plate
in aerospace components.
Sodium and its compounds
(1) Ores of sodium : NaCl (common salt),
3
NaNO
(chile salt petre), Na SO HO 2 4 2
. 10 (Glauber's salt), borax
(sodium tetraborate or sodium borate,
2 4 7 2
NaBO HO.
(2) Extraction of sodium : It is manufactured by
the electrolysis of fused sodium chloride in the
presence of 2
CaCl and KF using graphite anode and
iron cathode. This process is called Down process.
NaCl ⇌
Na Cl.
At cathode : Na e Na
At anode :
Cl Cl e ;
2
Cl Cl Cl
Sodium cannot be extracted from aqueous NaCl
because
0
/ 2 2
HOH
E
(–0.83 V ) is more than E Na / Na
0
2.71 V ).
Anode and cathode are separated by means of a
wire gauze to prevent the reaction between Na
and
2
Cl.
(3) Compound of sodium
(i) Sodium chloride : It is generally obtained by
evaporation of sea water by sun light. However NaCl so
obtained contains impurities like
4 2
CaSO , CaCl and
2
MgCl which makes the salt deliquescent. It is then
purified by allowing HCl
gas to pass through the
impure saturated solution of NaCl. The concentration
of
Cl ions increases and as a result pure NaCl gets
precipitated due to common ion effect.
(ii) Sodium hydroxide NaOH (Caustic soda)
Preparation
(a) Gossage process :
2 3 2 3
( 10 % solution)
Na CO CaOH NaOH CaCO
(b) Electrolytic method : Caustic soda is
manufactured by the electrolysis of a concentrated
solution of NaCl.
At anode:
Cl discharged; At cathode:
Na
discharged
(c) Castner - Kellener cell (Mercury cathode
process) : NaOH obtained by electrolysis of aq.
solution of brine. The cell comprises of rectangular iron
tank divided into three compartments.
Outer compartment – Brine solution is
electrolysed ; Central compartment – 2% NaOH
solution and
2
H
Properties : White crystalline solid, highly
soluble in water, It is only sparingly soluble in alcohol.
(a) Reaction with salt :
FeCl 3 NaOH
3
Fe ( OH ) 3 NaCl
(Insoluble hy droxide)
3
yellow
H
unstable
2 2 ( )
2 2 2
HgCl NaOH NaCl HgOH O HgO
AgNO 2 NaOH 2 NaNO 2 AgOH
3 3
Ag O HO
2
Brown
2
Zn , Al , Sb , Pb , Sn and As forms insoluble hydroxide
which dissolve in excess of NaOH (amphoteric
hydroxide).
NH Cl NaOH NaCl NH HO
3 2
heat
4
(b) Reaction with halogens :
X 2 NaOH
2
(cold)
sod. hypohalite
2
NaX NaXO HO
3 X 6 NaOH
2
(hot)
(Sod. halate)
3 2
5 NaX NaXO 3 HO ;
( X Cl , Br , I )
(c) Reaction with metals : Weakly electropositive
metals like Zn , Al and Sn etc.
2 2 2
Zn 2 NaOH NaZnO H
(d) Reaction with sand, SiO 2
2
2 NaOH SiO Na SiO HO
2
Sod. silicate(glass)
2 3
(e) Reaction with CO :
510 Sod.formate
150 200
NaOH CO HCOONa
atm
C
o
NaOH breaks down the proteins of the skin flesh
to a pasty mass, therefore it is commonly known as
caustic soda.
Caustic property : sodium hydroxide breaks
down the proteins of the skin flesh to a pasty mass,
therefore, it is commonly known as caustic soda.
Uses : Sodium hydroxide is used :
(a) in the manufacture of soidum metal, soap
(from oils and fats), rayon, paper, dyes and drugs,
(b) for mercurinzing cotton to make cloth
unshrinkable and
(c) as a reagent in the laboratory.
(iii) Sodium carbonate or washing soda,
2 3
NaCO
It exists in various forms, namely anhydrous sodium
carbonate Na 2 CO 2 (soda-ash); monohydrate Na CO HO 2 3 2
(crystal carbonate); hyptahydrate Na CO HO
2 3 2
. 7 and
decahydrate Na CO HO 2 3 2
. 10 (washing soda or sal soda).
Preparation : (a) Solvay process : In this process,
brine ( NaCl ),
3
NH
and
2
CO
are the raw materials.
NH CO HO
3 2 2 4 3
NHHCO
NaHCO NHCl
C
o
NH HCO NaCl
4 3 3 4
30
3 2 3 2 2
250
2 NaCO HO CO
C
o
NaHCO
4 2 2 2 3
lime
slaked
2 NH Cl Ca ( OH ) CaCl HO NH
2
CaCl
so formed in the above reaction is a by
product of solvay process.
Properties
(a)
Na CO HO NaCO HO HO
2 3 2 2 3 2 2
9
(M onohydrate)
.
dryair
(decahydrate)
. 10
.
2 3 2 2 3
Na CO HO NaCO
It does not decompose on funrther heating even to
redness (m.pt. 853° C )
(b) It is soluble in water with considerable
evolution of heat.
Na CO HO HCO 2 Na 2 OH
Weak acid
2 3 2 2 3
(c) It is readily decomposed by acids with the
evolution of
2
CO gas.
(d)
2 3 2 2 3
Na CO HO CO 2 NaHCO
Uses : In textile and petroleum refining,
Manufacturing of glass, NaOH soap powders etc.
(iv) Sodium peroxide (Na 2 O 2 )
Preparation : It is manufactured by heating
sodium metal on aluminium trays in air (free from )
2
CO
2
2 Na O (air)
2 2
NaO
Properties : (a) When pure it is colourless. The
faint yellow colour of commercial product is due to
presence of small amount of superoxide( ).
2
NaO
(b) On coming with moist air it become white due
to formation of NaOH and
2 3
NaCO.
2 2 2 2
2 Na O 2 HO 4 NaOH O ;
NaOH CO NaCO HO
2 2 3 2
(c) It is powerful oxidising agent. It oxidises
Cr (III) hydroxide to sodium chromate, Mn (II) to
sodium manganate and sulphides to sulphates.
Uses : As a bleaching agent and it is used for the
purification of air in confined spaces such as
submarines because it can combine with
2
CO to give
2 3
NaCO and oxygen,
2 2 2
2 CO 2 NaO
2 3 2
2 Na CO O.
(v) Micro cosmic salt [ Na ( NH 4 ) HPO 4. 4 H 2 O ]
Prepared by dissolving equimolar amounts of
2 4
NaHPO and NH Cl
4
in water in 1 : 1 ratio followed by
crystallization
Crystallization
NH Cl NaHPO NaNH HPO NaCl
4 2 4 4 4
(Colourles scry stal)
4 4 2
Na ( NH ) HPO. 4 HO
Chemical properties :
On heating M.C.S,
3
NaPO is formed.
3
NaPO forms
coloured beads with oxides of transition metal cloudy
2
SiO
2 3
phosphate)
(Sodium meta
4 4 3
Na ( NH ) HPO NaPO HO NH
(blue bead)
4
glassy bead)
(Trans parent
3
NaPO CuO CuNaPO
3 4
NaPO CoO CoNaPO (blue bend)
3 4
NaPO MnO NaMnO (blue bead)
Uses : (a) For the formation of sodium meta
phosphate and copper sodium phosphate
(b) It is used for the detection of colured ion
of the same period. This is due to the fact that alkaline
earth metals possess a higher nuclear charge than
alkali metals which more effectively pulls the orbit
electrons towards the nucleus causing a decrease in
size.
(3) Density
(i) Density decreases slightly upto Ca after which
it increases. The decrease in density from Be to Ca
might be due to less packing of atoms in solid lattice of
Mg and Ca.
Be Mg Ca Sr Ba Ra
(ii) The alkaline earth metals are more denser,
heavier and harder than alkali metal. The higher
density of alkaline earth metals is due to their smaller
atomic size and strong intermetallic bonds which
provide a more close packing in crystal lattice as
compared to alkali metals.
(4) Melting point and Boiling point
(i) Melting points and boiling points of alkaline
earth metals do not show any regular trend.
Be Mg Ca Sr Ba
Ra
melting points (K) 1560 920 1112 1041
1000 973
boiling point (K) 2770 1378 1767 1654 1413
(ii) The values are, however, more than alkali
metals. This might due to close packing of atoms in
crystal lattice in alkaline earth metals.
(5) Ionisation energy and electropositive or
metallic character
(i) Since the atomic size decreases along the
period and the nuclear charge increases and thus the
electrons are more tightly held towards nucleus. It is
therefore alkaline earth metals have higher ionisation
energy in comparison to alkali metals but lower
ionisation energies in comparison to p-block elements.
(ii) The ionisation energy of alkaline earth metals
decreases from Be to Ba.
Be Mg Ca Sr Ba
Ra
First ionisation energy ( k J mol
) 899 737 590 549
503 509
Second ionisation energy ( kJ mol
) 1757 1450 1146 1064
965 979
(iii) The higher values of second ionisation energy
is due to the fact that removal of one electron from the
valence shell, the remaining electrons are more tightly
held in which nucleus of cation and thus more energy is
required to pull one more electron from monovalent
cation.
(iv) No doubt first ionisation energy of alkaline
earth metals are higher than alkali metals but a closer
look on 2nd ionisation energy of alkali metals and
alkaline earth metals reveals that 2nd ionisation energy
of alkali metals are more
Li Be
1st ionisation energy ( kJ mol
) 520 899
2nd ionisation energy ( kJ mol
) 7296 1757
This may be explained as,
Li : 1s
2
, 2s
1
electron
removal of s
2
Li
: 1s
2
electron
removal ofs
1
Li
2+
: 1s
1
Be : 1s
2
, 2s
2
electron
removal of s
2
Be
: 1s
2
, 2s
1
electron
removal of s
2
Be
2+
:
1s
2
The removal of 2
nd
electron from alkali metals
takes place from 1s sub shell which are more closer to
nucleus and exert more nuclear charge to hold up 1s
electron core, whereas removal of 2nd electron from
alkaline earth metals takes from 2s sub shell. More
closer are shells to the nucleus, more tightly are held
electrons with nucleus and thus more energy is
required to remove the electron.
(v) All these possess strong electropositive
character which increases from Be to Ba.
(vi) These have less electropositive character than
alkali metals as the later have low values of ionisation
energy.
(6) Oxidation number and valency
(i) The IE 1 of the these metals are much lower
than IE 1 and thus it appears that these metals should
form univalent ion rather than divalent ions but in
actual practice, all these give bivalent ions. This is due
to the fact that M
2+
ion possesses a higher degree of
hydration or M
2+
ions are extensively hydrated to form
[ M ( H
2
O )
x
]
2+
, a hydrated ion. This involves a large
amount of energy evolution which counter balances the
higher value of second ionisation energy.
M M
2+
, H = IE
1
+ E
2
M
2+
- x H 2 O [ M ( H 2 O )x]
2+
; H = – hydration
energy.
(ii) The tendency of these metals to exist as
divalent cation can thus be accounted as,
(a) Divalent cation of these metals possess noble
gas or stable configuration.
(b) The formation of divalent cation lattice leads
to evolution of energy due to strong lattice structure of
divalent cation which easily compensates for the higher
values of second ionisation energy of these metals.
(c) The higher heats of hydration of divalent
cation which accounts for the existence of the divalent
ions of these metals in solution state.
(7) Hydration of ions
(i) The hydration energies of alkaline earth
metals divalent cation are much more than the
hydration energy of monovalent cation.
Mg
+
Mg
2+
Hydration energy or Heat of hydration ( kJ mol
) 353
1906
The abnormally higher values of heat of hydration
for divalent cations of alkaline earth metals are
responsible for their divalent nature. MgCl 2
formation
occurs with more amount of heat evolution and thus
MgCl 2
is more stable.
(ii) The hydration energies of M
2+
ion decreases
with increase in ionic radii.
Be
2+
Mg
2+
Ca
2+
Sr
2+
Ba
2+
Heat of hydration kJ mol
2382 1906 1651 1484
1275
(iii) Heat of hydration are larger than alkali
metals ions and thus alkaline earth metals compounds
are more extensively hydrated than those of alkali
metals e.g MgCl 2
and CaCl 2
exists as Mg Cl 2
.6 H
2
O and
CaCl 2. 6 H 2 O which NaCl and KCl do not form such
hydrates.
(iv) The ionic mobility, therefore, increases from
Be
2+
to Ba
2+
, as the size of hydrated ion decreases.
(8) Electronegativities
(i) The electronegativities of alkaline earth metals
are also small but are higher than alkali metals.
(ii) Electronegativity decreases from Be to Ba as
shown below,
Be Mg Ca Sr Ba
Electronegativity 1.57 1.31 1.00 0.
(9) Conduction power : Good conductor of heat
and electricity.
(10) Standard oxidation potential and reducing
properties
(i) The standard oxidation potential (in volts) are,
Be Mg Ca Sr Ba
1.69 2.35 2.87 2.89 2.
(ii) All these metals possess tendency to lose two
electrons to give M
2+
ion and are used as reducing
agent.
(iii) The reducing character increases from Be to
Ba , however, these are less powerful reducing agent
than alkali metals.
(iv) Beryllium having relatively lower oxidation
potential and thus does not liberate H 2 from acids.
(11) Characteristic flame colours
The characteristic flame colour shown are : Ca -
brick red; Sr – crimson ; Ba - apple green and Ra -
crimson.
Chemical Properties
(1) Formation of oxides and hydroxides
(i) The elements (except Ba and Ra ) when burnt
in air give oxides of ionic nature M
2+
O
2 -
which are
crystalline in nature. Ba and Ra however give peroxide.
The tendency to form higher oxides increases from Be
to Ra.
2 M + O 2 2 MO ( M is Be , Mg or Ca )
2 M + O
2
MO
2
( M is Ba or Sr )
(ii) Their less reactivity than the alkali metals is
evident by the fact that they are slowly oxidized on
exposure to air, However the reactivity of these metals
towards oxygen increases on moving down the group.
(iii) The oxides of these metals are very stable
due to high lattice energy.
(iv) The oxides of the metal (except BeO and MgO )
dissolve in water to form basic hydroxides and evolve a
large amount of heat. BeO and MgO possess high lattice
energy and thus insoluble in water.
(v) BeO dissolves both in acid and alkalies to give
salts i.e. BeO possesses amphoteric nature.
BeO+ 2 NaOH Na 2
BeO 2
+H
2
O ; BeO+ 2 HCl BeCl 2
+H
2
O
Sod. beryllate
Beryllium chloride
(vi)The basic nature of oxides of alkaline earth
metals increases from Be to Ra as the electropositive
Character increases from Be to Ra.
(vii)The tendency of these metal to react with
water increases with increase in electropositive
character i.e. Be to Ra.
(viii) Reaction of Be with water is not certain,
magnesium reacts only with hot water, while other
metals react with cold water but slowly and less
energetically than alkali metals.
(ix) The inertness of Be and Mg towards water is
due to the formation of protective, thin layer of
hydroxide on the surface of the metals.
(x) The basic nature of hydroxides increase from
Be to Ra. It is because of increase in ionic radius down
the group which results in a decrease in strength of M –
O bond in M – ( OH ) 2
to show more dissociation of
hydroxides and greater basic character.
(xi) The solubility of hydroxides of alkaline earth
metals is relatively less than their corresponding alkali
metal hydroxides Furthermore, the solubility of
hydroxides of alkaline earth metals increases from Be
to Ba. Be ( OH ) 2 and Mg ( OH ) 2 are almost insoluble, Ca
( OH )
2
(often called lime water) is sparingly soluble
and MgCl 2 the chlorides of alkaline earth metals impart
characteristic colours to flame.
CaCl 2
SrCl 2
BaCl 2
Brick red colour Crimson colour Grassy green
colour
Structure of BeCl 2
: In the solid phase polymeric
chain structure with three centre two electron bonding
with Be - C l - Be bridged structure is shown below,
In the vapour phase it tends to form a chloro-
bridged dimer which dissociates into the linear
triatomic monomer at high temperature at nearly 1200
K.
(5) Solubility in liquid ammonia : Like alkali
metals, alkaline earth metals also dissolve in liquid
ammonia to form coloured solutions When such a
solution is evaporated, hexammoniate, M ( NH 3 ) 6 is
formed.
(6) Nitrides
(i) All the alkaline earth metals direct combine
with N 2
give nitrides, M 3
N
2
(ii) The ease of formation of nitrides however
decreases from Be to Ba.
(iii) These nitrides are hydrolysed water to
liberate
NH
3
, M
3
N
2
+ 6 H
2
O 3 M ( OH )
2
+ 2 NH
3
(7) Sulphates
(i) All these form sulphate of the type M SO 4
by
the action of H 2 SO 4 on metals, their oxides, carbonates
or hydroxides.
M + H
2
SO
4
MSO
4
+ H
2
; MO + H
2
SO
4
MSO 4 + H 2 O
MCO
3
+ H
2
SO
4
MSO
4
+ H
2
O + CO
2
M ( OH )
2
+ H
2
SO
4
MSO
4
+ 2 H
2
O
(ii) The solubility of sulphates in water decreases
on moving down the group BeSO 4 and MgSO 4 are fairly
soluble in water while BaSO 4 is completely insoluble.
This is due to increases in lattice energy of sulphates
down the group which predominates over hydration
energy.
(iii) Sulphate are quite stable to heat however
reduced to sulphide on heating with carbon.
MSO
4
+ 2 C MS +2 CO
2
(8) Action with carbon : Alkaline metals (except
Be , Mg ) when heated with carbon form carbides of the
type MC 2 These carbides are also called acetylides as on
hydrolysis they evolve acetylene.
MC 2 + 2 H 2 O M ( OH ) 2 + C 2 H 2
(9) Action with sulphur and phosphorus :
Alkaline earth metals directly combine with sulphur
and phosphorus when heated to form sulphides of the
type MS and phosphides of the type M 3 P 2 respectively.
M + S MS ; 3 M + 2 P M
3
P
2
Sulphides on hydrolysis liberate H 2 S while
phosphides on hydrolysis evolve phosphine.
MS + dil. acid H 2
S ; M
3
P
2
Sulphides are phosphorescent and are
decomposed by water
2 MS + 2H
2
O M ( OH )
2
+ M ( HS )
2
(10) Nitrates : Nitrates of these metals are soluble
in water On heating they decompose into their
corresponding oxides with evolution of a mixture of
nitrogen dioxide and oxygen.
3 2 2 2
2
1
M ( NO ) MO 2 NO O
(11) Formation of complexes
(i) Tendency to show complex ion formation
depends upon smaller size, high nuclear charge and
vacant orbitals to accept electron. Since alkaline metals
too do not possess these characteristics and thus are
unable to form complex ion.
(ii) However, Be
2+
on account of smaller size
forms many complex such as ( BeF 3 )
1 -
, ( BeF 4 )
2 -
Anomalous behaviour of Beryllium
Beryllium differs from rest of the alkaline earth
metals on account of its small atomic size, high
electronegativity Be
2+
exerts high polarizing effect on
anions and thus produces covalent nature in its
compounds. Following are some noteworthy difference
of Be from other alkaline earth metals,
(1) Be is lightest alkaline earth metal.
(2) Be possesses higher m.pt. and b.pt than other
group members.
(3) BeO is amphoteric in nature whereas oxides of
other group members are strong base.
(4) It is not easily effected by dry air and does not
decompose water at ordinary temperature.
(5) BeSO 4 is soluble in water.
(6) Be and Mg carbonates are not precipitated by
4 2 3
( NH ) CO in presence of NH 4 Cl.
(7) Be and Mg salts do not impart colour to flame.
(8) Be does not form peroxide like other alkaline
earth metals.
(9) It does not evolve hydrogen so readily from
acids as other alkaline earth metals do so.
(10) It has strong tendency to form complex
compounds.
(11) Be 3
N
2
is volatile whereas nitrides of other
alkaline earth metals are non-volatile.
Cl
Cl
Be
263
pm
Cl
Cl
Be
98
o
Cl
Cl
Be
202 PM
82
o
(12) It’s salts can never have more than four
molecules of water of crystallization as it has only four
available orbitals in its valence shell.
(13) Berylium carbide reacts water to give
methane whereas magnesium carbide and calcium
carbide give propyne and acetylene respectively.
Be 2
C +4 H
2
O 2 Be ( OH ) 2
+ CH
4
Mg 2
C
3
+ 4 H
2
O 2 Mg ( OH ) 2
+ C
3
H
6
CaC 2
+ 2 H
2
O Ca ( OH ) 2
+ C
2
H
4
Diagonal relationship of Be with Al
Due to its small size Be differs from other earth
alkaline earth metals but resembles in many of its
properties with Al on account of diagonal relationship.
(1) Be
2+
and Al
3+
have almost same and smaller
size and thus favour for covalent bonding.
(2) Both these form covalent compounds having
low m. pt and soluble in organic solvent.
(3) Both have same value of electronegativity ( i.e.
(4) The standard O.P of these elements are quite
close to each other ; Be
2+
=1.69 volts and Al
3+
volts.
(5) Both become passive on treating with conc.
HNO
3
in cold.
(6) Both form many stable complexes e.g. ( BeF 3
( AlH 4 )
(7) Like BeO , Al 2 O 3 is amphoteric in nature. Also
both are high melting point solids.
Al 2 O 3 + 2 NaOH 2 NaAlO 2 + H 2 O
Al 2 O 3 + 6 HCl 2 AlCl 3 + 3 H 2 O
(8) Be and Al both react with NaOH to liberate H 2
forming beryllates and alluminates.
Be + 2 NaOH Na 2 BeO 2 + H 2
2 Al + 6 NaOH 2 Na 3 AlO 3 + 3 H 2
(9) Be 2 C and Al 4 C 3 both give CH 4 on treating with
water.
Be 2
C + 2 H
2
O CH
4
Al 4 C 3 + 6 H 2 O 3 CH 4 + 2 Al 2 O 3
(10) Both occur together in nature in beryl ore,
3 BeO. Al 2
O
6 SiO 2
(11) Unlike other alkaline earths but like
aluminium, beryllium is not easily attacked by air (Also
Mg is not attacked by air)
(12) Both Be and Al react very slowly with dil. HCl
to liberate H 2
(13) Both Be and Al form polymeric covalent
hydrides while hydrides of other alkaline earth are
ionic.
(14) Both BeCl 2 and AlCl 3 are prepared is similar
way.
BeO + C + Cl 2 BeCl 2 + CO
Al 2 O 3 + 3 C +3 Cl 2 2 AlCl 3 + 3 CO
(15) Both BeCl 2
and AlCl 3
are soluble in organic
solvents and act as catalyst in Friedel – Crafts reaction.
(16) Both Be ( OH ) 2 and Al ( OH ) 3 are amphoteric
whereas hydroxides of other alkaline earths are strong
alkali.
(17) The salts of Be and Al are extensively
hydrated.
(18) BeCl 2
and AlCl 3
both have a bridged polymeric
structure.
(19) Be and Al both form fluoro complex ions
[ BeF 4 ]
2 –
and [ AlF 6 ]
3 –
in solution state whereas other
members of 2nd group do not form such complexes.
Magnesium and its compounds
(1) Ores of magnesium : Magnesite ( ),
3
MgCO
Dolomite (. )
3 3
MgCO CaCO , Epsomite (epsom salt)
4 2
MgSO HO Carnallite (.. 6 )
2 2
MgCl KCl HO Asbestos
3 34
CaMg SiO Talc
2
( Mg ( ). ( ))
2 52 2
SiO MgOH.
(2) Extraction of magnesium : It is prepared by
the electrolysis of fused magnesium chloride which in
turn is obtained from carnallite and magnesite.
Carnallite (.. 6 )
2 2
MgCl KCl HO can’t be directly
converted into anhydrous
2
MgCl by heating because all
the water of crystallisation cannot be removed by
heating. Moreover, strong heating may change it to
MgO.
MgCl HO MgO HCl HO
2 2 2
2 2
In Dow’s process, magnesium chloride is obtained
from sea water as MgCl HO
2 2
. It is rendered
anhydrous by heating it in a current of dry HCl gas.
The anhydrous magnesium chloride is fused with
NaCl (to provide conductivity to the electrolyte and to
lower the fusing temperature of anhydrous
2
MgCl ) and
then electrolysed at C
o
(3) Compounds of magnesium
(i) Magnesia ( MgO ) : It is used as magnesia
cement. It is a mixture of
MgO and.
2
MgCl It is also
called Sorel's cement.
(ii) Magnesium hydroxide : It aqueous
suspension is used in Medicine as an antacid. Its
medicinal name is milk of magnesia.
(iii) Magnesium sulphate or Epsom salt
4 2
MgSO HO : It is isomorphous with. 7.
4 2
ZnSO HO It is
used as a purgative in medicine, as a mordant in dyeing
and as a stimulant to increase the secretion of bile.
(iv) Magnesium chloride (. 6 )
2 2
MgCl HO : It is a
deliquescent solid. Hydrated salt on heating in air
undergoes partial hydrolysis.
MgCl HO MgOHCl HCl HO
2 2 2
( ) 5
Heat
Calcium and its compounds
Mendeleefs periodic table) includes boron ( B ) ;
aluminium ( Al ) , gallium ( Ga ), indium ( In ) and thallium
( Tl ) Boron is the first member of group 13 of the
periodic table and is the only non-metal of this group.
The all other members are metals. The non-metallic
nature of boron is due its small size and high ionisation
energy. The members of this family are collectively
known as boron family and sometimes as aluminium
family.
(1) Electronic configuration
Element
Electronic configuration
2 1
ns np )
B
5
2 1
[ He ] 2 s 2 p
Al
13
2 1
[ Ne ] 3 s 3 p
Ga
31
10 2 1
[ Ar ] 3 d 4 s 4 p
In
49
10 2 1
[ Kr ] 4 d 5 s 5 p
Tl
81
14 10 2 1
[ Xe ] 4 f 5 d 6 s 6 p
(2) Occurrence : The important of this group
elements are given below,
Boron : Borax (Tincal) ( Na 2 B 4 O 7 .10 H 2 O ), Colemanite
( Ca 2
B
6
O
11
5 H
2
O )
Boracite (2 Mg 3 B 8 O 15. MgCl 2 ), Boronatro calcite
( CaB 4 O 7. NaBO 2 .8 H 2 O ),
Kernite ( Na 2 B 4 O 7. 4 H 2 O ), Boric acid ( H 3 BO 3 )
Aluminium : Corundum ( Al 2
O
3
), Diaspore
(Al 2
O
3
. H
2
O ), Bauxite ( Al 2
O
2 H
2
O ), and Cryolite
( Na 3
AlF 6
Physical properties
(1) A regular increasing trend in density down the
group is due to increase in size.
(2) Melting points do not vary regularly and
decrease from B to Ga and then increase.
(3) Boron has very high melting point because it
exist as giant covalent polymer in both solid and liquid
state.
(4) Low melting point of Ga (29.
0
C ) is due to the
fact that consists of only Ga 2 molecule; it exist as liquid
upto 2000
0
C and hence used in high temperature
thermometry.
(5) Boiling point of these elements however show
a regular decrease down the group.
(6) The abrupt increase in the atomic radius of Al
is due to greater screening effect in Al (it has 8
electrons in its penultimate shell) than in B (it has 2
electrons in its penultimate shell)
(7) The atomic radii of group 13 elements are
smaller than the corresponding s-block elements. This
is due to the fact that when we move along the period,
the new incoming electron occupy the same shell
whereas the nuclear charge increases regularly
showing more effective pull of nucleus towards shell
electrons. This ultimately reduces the atomic size.
(8) The atomic radius of Ga is slightly lesser than
of Al because in going from Al to Ga , the electrons have
already occupied 3 d sub shell in Ga. The screening
effect of these intervening electrons being poor and has
less influence to decrease the effective nuclear charge,
therefore the electrons in Ga experience more forces of
attractions towards nucleus to result in lower size of
Ga than Al
(9) Oxidation state
(i) All exhibit +3 oxidation state and thus
complete their octet either by covalent or ionic union.
(ii) Boron being smaller in size cannot lose its
valence electrons to form B
3+
ion and it usually show +
covalence. The tendency to show +3 covalence however
decreases down the group even Al shows +3 covalence
in most of its compounds.
(iii) Lower elements also show +1 ionic state e.g
Tl
, Ga
. This is due to inert pair effect. The
phenomenon in which outer shell ‘ s ’ electrons ( ns
2
penetrate to ( n - 1) d - electrons and thus become closer
to nucleus and are more effectively pulled the nucleus.
This results in less availability of ns
2
electrons pair for
bonding or ns
2
electron pair becomes inert. The inert
pair effect begins after n 4 and increases with
increasing value of n.
(iv) The tendency to form M
ion increases down
the gp. Ga
< Tl
(10) Hydrated ions : All metal ions exist in
hydrated state.
(11) Ionisation energy
(i) Inspite of the more charge in nucleus and
small size, the first ionisation energies of this group
elements are lesser than the corresponding elements of
s block. This is due to the fact that removal of electron
from a p - orbitals (being far away from nucleus and
thus less effectively held than s-orbitals) is relatively
easier than s - orbitals.
(ii) The ionisation energy of this group element
decrease down the group due to increases in size like
other group elements.
(iii) However, ionisation energy of Ga are higher
than that of Al because of smaller atomic size of Ga due
to less effective shielding of 3 d electrons in Ga. Thus
valence shell exert more effective nuclear charge in Ga
to show higher ionisation energies.
(12) Electropositive character
(i) Electropositive character increases from B to
Tl.
(ii) Boron is semi metal, more closer to non-
metallic nature whereas rest all members are pure
metals.
(iii) Furthermore, these elements are less
electropositive than s - block elements because of
smaller size and higher ionisation energies.
(13) Oxidation potential
(i) The standard oxidation potentials of these
element are quite high and are given below,
B Al Ga In Tl
E
0
op for M M
3+
+1.
E
0
op for M M
+0.
(ii) However Boron does not form positive ions in
aqueous solution and has very low oxidation potential.
(iii) The higher values of standard oxidation
potentials are due to higher heats of hydration on
account of smaller size of trivalent cations.
(iv) Aluminium is a strong reducing agent and can
reduce oxides which are not reduced even by carbon.
This is due to lower ionisation energy of aluminium than
carbon. The reducing character of these elements is A l > Ga
In > Tl.
(14) Complex formation : On account of their
smaller size and more effective nuclear charge as well
as vacant orbitals to accept elements, these elements
have more tendency to form complexes than- s block
elements.
Chemical properties
(1) Hydrides
(i) Elements of group 13 do not react directly with
hydrogen but a number of polymeric hydrides are
known to exist.
(ii) Boron forms a large no. of volatile covalent
hydrides, known as boranes e.g. B 2 H 6 , B 4 H 10 , B 5 H 11 , B 6 H 10
Two series of borones with general formula B n
H
n+
and
Bn H n+6 are more important.
(iii) Boranes are electron deficient compounds. It
is important to note that although BX 3
are well known,
BH 3 is not known. This is due of the fact that hydrogen
atoms in BH 3 have no free electrons to form p – p back
bonding and thus boron has incomplete octet and
hence BH 3 molecules dimerise to form B 6 H 6 having
covalent and three centre bonds.
(iv) Al forms only one polymeric hydride ( AlH 3 )n
commonly known as alane It contains Al ….. H …… Al
bridges.
(v) Al and Ga forms anionic hydrides e.g. LiAlH 4 and
LiGa H 4
LiH AlCl LiAlH LiCl
ether
4 [ ] 3
3 4
(2) Reactivity towards air
(i) Pure boron is almost unreactive at ordinary
temperature. It reacts with air to form B 2 O 3 when
heated It does react with water. Al burns in air with
evolution of heat give Al 2
O
3
(ii) Ga and In are not effected by air even
when heated whereas Tl is little more reactive and also
form an oxide film at surface. In moist air, a layer of Tl
( OH ) is formed.
(iii) Al decomposes H 2
O and reacts readily in air
at ordinary temperature to form a protective film of its
oxides which protects it from further action.
(3) Oxides and hydroxides
(i) The members of boron family form oxide and
hydroxides of the general formula M 2 O 3 and M ( OH ) 3
respectively.
(ii) The acidic nature of oxides and hydroxides
changes from acidic to basic through amphoteric from
B to Tl.
B 2 O 3 and B ( OH ) 3 > Al 2 O 3 and Al ( OH ) 3 >
(acidic) (amphoteric)
Ga 2 O 3 and Ga ( OH ) 3 > In 2 O 3 In ( OH ) 3 > Tl 2 O 3
Tl ( OH ) 3
(amphoteric) (basic) (strong
basic)
B ( OH )
3
or H 3
BO
3
is weak monobasic Lewis acid.
(iii) Boric acid, B ( OH ) 3
is soluble in water as it
accepts lone pair of electron to act as Lewis acid. Rest
all hydroxides of group 13 are insoluble in water and
form a gelatinous precipitate.
B ( OH )
3
+ H
2
O B ( OH )
4
1 –
+ H
(iv) Al 2
O
3
being amphoteric dissolves in acid and
alkalies both.
Al 2
O
3
+ 3 H
2
SO
4
Al 2
( SO
4
3
+ 3 H
2
O
Al O NaOH NaAlO HO
fuse
2
Sodium metaaluminate
2 3 3
2 2
(v) One of the crystalline form of alumina ( Al 2
O
3
is called corrundum. It is very hard and used as
abrasive. It is prepared by heating amorphous form of
Al 2
O
3
to 2000 K.
(4) Action of Acids
(i) Boron does not react with non oxidizing acids,
however, it dissolves in nitric acid to form boric acids.
(ii) Al, Ga and In dissolve in acids forming their
trivalent cations; however, Al and Ga become passive
due to the formation of protective film of oxides.
(iii) Thallium dissolves in acids forming univalent
cation and becomes passive in HCl due to the formation
of water insoluble TICl.
into [ M ( H 2 O ) 6 ]
3+
and 3 X
ions and the solution becomes
good conductor of electricity.
Al 2 Cl 6 + 2 H 2 O 2[ Al ( H 2 O ) 6 ]
3+
+6 Cl
; Therefore
Al 2
Cl 6
is ionic in water.
The dimeric structure may also split by reaction
with donor molecules e.g. R 3 N. This is due to the
formation of complexes of the type R 3
NAlCl 3
The
dimeric structure of Al 2 Cl 6 exist in vapour state below
473 K and at higher temperature it dissociates to
trigonal planar AlCl 3
molecule.
Boron halides do not exist as dimer due to small
size of boron atom which makes it unable to co-
ordinate four large-sized halide ions.
(x) BF 3
and AlCl 3
acts as catalyst and Lewis acid in
many of the industrial process.
Anomalous Behaviour of Boron
Like Li and Be , Boron – the first member of group
13 also shows anomalous behaviour due to extremely
low size and high nuclear charge/size ratio, high
electronegativity and non-availability of d electrons.
The main point of differences are,
(1) Boron is a typical non- metal whereas other
members are metals.
(2) Boron is a bad conductor of electricity
whereas other metals are good conductors.
(3) Boron shows allotropy and exists in two forms
- crystalline and amorphous. Aluminium is a soft metal
and does not exist in different forms.
(4) Like other non-metals, the melting point and
boiling point of boron are much higher than those of
other elements of group 13.
(5) Boron forms only covalent compounds
whereas aluminium and other elements of group 13
form even some ionic compounds.
(6) The hydroxides and oxides of boron are acidic
in nature whereas those of others are amphoteric and
basic.
(7) The trihalides of boron ( BX 3 ) exist as monomers
On the other hand, aluminium halides exist as dimers
( Al 2
X
6
(8) The hydrides of boron i.e. boranes are quite
stable while those of aluminium are unstable.
(9) Dilute acids have no action on boron Others
liberate H 2 from them.
(10) Borates are more stable than aluminates.
(11) Boron exhibit maximum covalency of four
e.g., BH
4
ion while other members exhibit a maximum
covalency of six e.g., [ Al ( OH ) 6 ]
3 -
(12) Boron does not decompose steam while other
members do so.
(13) Boron combines with metals to give borides
e.g. Mg 3
B
2
. Other members form simply alloys.
(14) Concentrated nitric acid oxidises boron to
boric acid but no such action is noticed other group
members.
B + 3 HNO 3 H 3 BO 3 + 3 NO 2
Diagonal relationship between Boron and
Silicon
Due to its small size and similar charge/mass
ratio, boron differs from other group 13 members, but
it resembles closely with silicon, the second element of
group 14 to exhibit diagonal relationship. Some
important similarities between boron and silicon are
given below,
(1) Both boron and silicon are typical non-metals,
having high m.pt. b.pt nearly same densities
( B =2.35 gml
S =2.34 g//ml ). low atomic volumes and
bad conductor of current. However both are used as
semiconductors.
(2) Both of them do not form cation and form only
covalent compounds.
(3) Both exists in amorphous and crystalline state
and exhibit allotropy.
(4) Both possess closer electronegativity values
( B =2.0; Si =1.8).
(5) Both form numerous volatile hydrides which
spontaneously catch fire on exposure to air and are
easily hydrolysed.
(6) The chlorides of both are liquid, fume in most
air and readily hydrolysed by water.
BCl 3 + 3 H 2 O B ( OH ) 3 + 3 HCl
SiCl 4
+ H
2
O Si ( OH ) 4
(7) Both form weak acids like H 3 BO 3 and H 2 SiO 3.
(8) Both form binary compounds with several
metals to give borides and silicide. These borides and
silicide react with H 3 PO 4 to give mixture of boranes and
silanes.
3 Mg +2 B Mg 3
B
2
; Mg 3
B
2
+ H
3
PO
4
Mixture of
boranes
(Magnesium boride)
2 Mg + Si Mg 2 Si ; Mg 2 Si + H 3 PO 4 Mixture of
silanes
(magnesium silicide)
(9) The carbides of both Boron and silicon ( B 4
C
and SiC ) are very hard and used as abrasive.
(10) Oxides of both are acidic and can be reduced
by limited amount of Mg In excess of Mg boride and
silicide are formed.
B 2 O 3 +3 Mg 3 MgO +2 B ; SiO 2 +2 Mg 2 MgO + Si
(11) Both the metals and their oxides are readily
soluble in alkalies.
2B + 6 NaOH 2 Na 3 BO 3 + 3 H 2
(borate)
Si + 2 NaOH + H 2 O Na 2 SiO 3 + 2 H 2
(silicate)
B
2
O
3
BO
3
+ 3 H
2
O
SiO 2 + 2 NaOH Na 2 SiO 3 + H 2 O
Both borates and silicates have tetrahedral
structural units
n
BO
4
and
n
SiO
4
respectively. Boro
silicates are known in which boron replaces silicon in
the three dimensional lattice. Boron can however form
planar BO 3
units.
(12) Acids of both these elements form volatile
esters on heating with alcohol in presence of conc.
H
2
SO
4
B ( OH ) 3 + 3 ROH B ( OR ) 3 + 3 H 2 O
Si ( OH ) 4
+ 4 H
2
O
Boron and its compounds
Boron is the first member of group – 13 (IIIA) of
the periodic table. Boron is a non- metal. It has a small
size and high ionization energy due to which it can not
lose its valence electrons to form
3
B ion. Its
compounds especially the hydrides and halides are
electron deficient and behave as Lewis acid.
(1) Ores of boron
(i) Borax or tincal : Na 2 B 4 O 7. 10 H 2 O
(ii) Kernite or Rasorite : Na 2
B
4
O
7
. 4 H
2
O
(iii) Colemanite : Ca 2 B 6 O 11. 5 H 2 O
(iv) Orthoboric acid : H 3 BO 3 (It occurs in the jets
of steam called soffioni escaping from ground in the
volcanic region of the Tuscany ). Boron is present to a
very small extent (0.001%) in earth’s crust.
(2) Isolation : Elemental boron in the form of
dark brown powder is obtained either by reduction of
boric oxide with highly electropositive metals like K ,
Mg , Al , Na , etc. in the absence of air and boron halides
with hydrogen at high temperature eg.
B
2
O
3
+ 6 K
Heat
2 B + 3 K
2
O
2 BCl 3
+ 3 H
1270 K
2 B + 6 HCl.
By thermal decomposition of boron triiodide over
red hot tungsten filament and boron hydrides for
example,
2 BI
3
W , heat
2 B + 3 I
2
; B
2
H
6
Heat
2B + 3 H
2
(3) Properties : It exists in mainly two allotropic
forms i.e. amorphous dark brown powder and
crystalline black very hard solid. It occurs in two
isotopic forms, i.e. ,
10
5
B (20% abundance) and
11
5
B
(80% abundance). With air, boron forms 2 3
BO and BN
at 973 K , with halogens, trihalides ( ) 3
BX are formed,
with metals borides are formed. eg.
4 B +
2
3 O
Heat
Boron trioxide
2 3
2 BO
2 B +
2
N
Heat
Boronnitride
2 BN
2 B +
2
3 X
Boron trihalide
3
2 BX
3 Mg + 2 B
Heat
Magnesium boride
3 2
MgB
Water, steam and HCl have no action on B.
oxidising acids
3
( HNO , )
2 4
HSO convert boron to
3 3
HBO
B + 3
3
HNO
3 3
HBO
2
3 NO
2 B + 3
2 4
HSO 2
3 3
HBO + 3
2
SO
Fused alkalies ( NaOH , KOH ) dissolve boron
forming borates, liberating hydrogen.
2 B + 6 KOH
Fused
3 3
KBO +
2
3 H
(4) Uses of Boron : Boron is used in atomic
reactors as protective shields and control rods, as a
semiconductors for making electronic devices in steel
industry for increasing the hardness of steel and in
making light composite materials for air crafts.
(5) Compounds of Boron
(i) Boron Hydrides
Boron forms hydrides of the types
n n 4
BH and
n n 6
BH
called boranes. Diborane is the simplest boron
hydride which is a dimer of
3
BH.
Preparation
(a)
2 6 4
450
3
8 BF 6 LiH BH 6 LiBF
K
(b) 4 BCl LiAlH 2 BH 3 AlCl 3 LiCl
3 4 2 6 3
(c) In the laboratory, it is prepared by the
oxidation of sod. Borohydride with
2
I.
2 6 2
Poly ether
4 2
2 NaBH I BH 2 NaI H
Properties : (a) Since Boron in boranes never
complete its octet of electrons hence all boranes are
called as electron-deficient compounds or Lewis acids.
(b) All boranes catch fire in the presence of
oxygen to liberated a lot of heat energy. Thus, they can
also be used as high energy fuels.
B H 3 O 2 BO 3 HO ; H 1976 KJ / mole
2 6 2 2 3 2
(c) Boranes are readily hydrolysed by water.
2 6 2 3 3 2
B H 6 HO 2 HBO 6 H
(d) With carbon monoxide
2 6 3 2
B H 2 CO ( BH CO )
(e) Boranes are used for formation of
hydroborates or borohydrides such as
4
LiBH
or
4
NaBH
which are extensively used as reducing agents in
organic synthesis.
2 2 [ ]
4
Diethy lether
2 6
LiH BH Li BH
2 2 [ ]
4
Diethy lether
2 6
NaH BH Na BH
Structure of diborane :
2 6
BH has a three centre
electon pair bond also called a banana shape bond.
(a)
t
B H
: It is a normal covalent bond (two
centre electron pair bond i.e. , 2 c 2 e ).
