315THE p-BLOCK ELEMENTS
UNIT 11
After studying this unit, you will be
able to
••
appreciate the general trends in the
chemistry of p-block elements;
••
••
describe the trends in physical and
chemical properties of group 13 and
14 elements;
••
••
explain anomalous behaviour of
boron and carbon;
••
••
describe allotropic forms of carbon;
••
••
know the chemistry of some
important compounds of boron,
carbon and silicon;
••
••
list the important uses of group 13
and 14 elements and their
compounds.
THE p -BLOCK ELEMENTS
In p-block elements the last electron enters the outermost
p orbital. As we know that the number of p orbitals is three
and, therefore, the maximum number of electrons that can
be accommodated in a set of p orbitals is six. Consequently
there are six groups of p–block elements in the periodic
table numbering from 13 to 18. Boron, carbon, nitrogen,
oxygen, fluorine and helium head the groups. Their valence
shell electronic configuration is ns
2
np
1-6
(except for He).
The inner core of the electronic configuration may,
however, differ. The difference in inner core of elements
greatly influences their physical properties (such as atomic
and ionic radii, ionisation enthalpy, etc.) as well as chemical
properties. Consequently, a lot of variation in properties of
elements in a group of p-block is observed. The maximum
oxidation state shown by a p-block element is equal to the
total number of valence electrons (i.e., the sum of the s-
and p-electrons). Clearly, the number of possible oxidation
states increases towards the right of the periodic table. In
addition to this so called group oxidation state, p-block
elements may show other oxidation states which normally,
but not necessarily, differ from the total number of valence
electrons by unit of two. The important oxidation states
exhibited by p-block elements are shown in Table 11.1. In
boron, carbon and nitrogen families the group oxidation
state is the most stable state for the lighter elements in the
group. However, the oxidation state two unit less than the
group oxidation state becomes progressively more stable
for the heavier elements in each group. The occurrence of
oxidation states two unit less than the group oxidation
states are sometime attributed to the ‘inert pair effect’.
The variation in properties of the p-block elements due to the
influence of d and f electrons in the inner core of the heavier
elements makes their chemistry interesting
2019-20
316 CHEMISTRY
The relative stabilities of these two oxidation
states group oxidation state and two unit less
than the group oxidation state may vary from
group to group and will be discussed at
appropriate places.
It is interesting to note that the non-metals
and metalloids exist only in the p-block of the
periodic table. The non-metallic character of
elements decreases down the group. In fact the
heaviest element in each p-block group is the
most metallic in nature. This change from non-
metallic to metallic character brings diversity
in the chemistry of these elements depending
on the group to which they belong.
In general, non-metals have higher ionisation
enthalpies and higher electronegativities than
the metals. Hence, in contrast to metals which
readily form cations, non-metals readily form
anions. The compounds formed by highly
reactive non-metals with highly reactive metals
are generally ionic because of large differences
in their electronegativities. On the other hand,
compounds formed between non-metals
themselves are largely covalent in character
because of small differences in their
electronegativities. The change of non-metallic
to metallic character can be best illustrated by
the nature of oxides they form. The non-metal
oxides are acidic or neutral whereas metal
oxides are basic in nature.
The first member of p-block differs from the
remaining members of their corresponding
group in two major respects. First is the size
and all other properties which depend on size.
Thus, the lightest p-block elements show the
same kind of differences as the lightest s-block
elements, lithium and beryllium. The second
important difference, which applies only to the
p-block elements, arises from the effect of d-
orbitals in the valence shell of heavier elements
(starting from the third period onwards) and
their lack in second period elements. The
second period elements of p-groups starting
from boron are restricted to a maximum
covalence of four (using 2s and three 2p
orbitals). In contrast, the third period elements
of p-groups with the electronic configuration
3s
2
3p
n
have the vacant 3d orbitals lying
between the 3p and the 4s levels of energy.
Using these d-orbitals the third period
elements can expand their covalence above
four. For example, while boron forms only
[BF
4
]
, aluminium gives [AlF
6
]
3–
ion. The
presence of these d-orbitals influences the
chemistry of the heavier elements in a number
of other ways. The combined effect of size and
availability of d orbitals considerably
influences the ability of these elements to form
π
bonds. The first member of a group differs
from the heavier members in its ability to form
p
π
- p
π
multiple bonds to itself ( e.g., C=C, CC,
Table 11.1 General Electronic Configuration and Oxidation States of p-Block Elements
Group 13 14 15 16 17 18
General
electronic ns
2
np
1
ns
2
np
2
ns
2
np
3
ns
2
np
4
ns
2
np
5
ns
2
np
6
configuration (1s
2
for He)
First member
of the B C N O F He
group
Group
oxidation +3 +4 +5 +6 +7 +8
state
Other
oxidation +1 +2, 4 +3, 3 +4, +2, –2 +5, + 3, +1, –1 +6, +4, +2
states
2019-20
317THE p-BLOCK ELEMENTS
NN) and to other second row elements (e.g.,
C=O, C=N, CN, N=O). This type of
π
- bonding
is not particularly strong for the heavier
p-block elements. The heavier elements do form
π
bonds but this involves d orbitals (d
π
– p
π
or d
π
–d
π
). As the d orbitals are of higher
energy than the p orbitals, they contribute less
to the overall stability of molecules than does
p
π
- p
π
bonding of the second row elements.
However, the coordination number in species
of heavier elements may be higher than for
the first element in the same oxidation state.
For example, in +5 oxidation state both N and
P form oxoanions : NO
3
(three-coordination
with
π
bond involving one nitrogen p-orbital)
and
3
4
PO
(four-coordination involving s, p and
d orbitals contributing to the
π
bond). In
this unit we will study the chemistry of group
13 and 14 elements of the periodic table.
11.1 GROUP 13 ELEMENTS: THE BORON
FAMILY
This group elements show a wide variation in
properties. Boron is a typical non-metal,
aluminium is a metal but shows many
chemical similarities to boron, and gallium,
indium, thallium and nihonium are almost
exclusively metallic in character.
Boron is a fairly rare element, mainly
occurs as orthoboric acid, (H
3
BO
3
), borax,
Na
2
B
4
O
7
·10H
2
O, and kernite, Na
2
B
4
O
7
·4H
2
O.
In India borax occurs in Puga Valley (Ladakh)
and Sambhar Lake (Rajasthan). The
abundance of boron in earth crust is less than
0.0001% by mass. There are two isotopic
forms of boron
10
B (19%) and
11
B (81%).
Aluminium is the most abundant metal and
the third most abundant element in the earth’s
crust (8.3% by mass) after oxygen (45.5%) and
Si (27.7%). Bauxite, Al
2
O
3
. 2H
2
O and cryolite,
Na
3
AlF
6
are the important minerals of
aluminium. In India it is found as mica in
Madhya Pradesh, Karnataka, Orissa and
Jammu. Gallium, indium and thallium are less
abundant elements in nature. Nihonium has
symbol Nh, atomic number 113, atomic mass
286 g mol
-1
and electronic configuration [Rn]
5f
14
6d
10
7s
2
7p
2
.
So far it has been prepared
in small amount and half life of its most stable
isotope is 20 seconds. Due to these reasons its
chemistry has not been established.
Nihonium is a synthetically prepared
radioactive element. Here atomic, physical and
chemical properties of elements of this group
leaving nihonium are discussed below.
11.1.1 Electronic Configuration
The outer electronic configuration of these
elements is ns
2
np
1
. A close look at the
electronic configuration suggests that while
boron and aluminium have noble gas
core, gallium and indium have noble gas plus
10 d-electrons, and thallium has noble gas
plus 14 f- electrons plus 10 d-electron cores.
Thus, the electronic structures of these
elements are more complex than for the first
two groups of elements discussed in unit 10.
This difference in electronic structures affects
the other properties and consequently the
chemistry of all the elements of this group.
11.1.2 Atomic Radii
On moving down the group, for each successive
member one extra shell of electrons is added
and, therefore, atomic radius is expected to
increase. However, a deviation can be seen.
Atomic radius of Ga is less than that of Al. This
can be understood from the variation in the
inner core of the electronic configuration. The
presence of additional 10 d-electrons offer
only poor screening effect (Unit 2) for the outer
electrons from the increased nuclear charge in
gallium. Consequently, the atomic radius of
gallium (135 pm) is less than that of
aluminium (143 pm).
11.1.3 Ionization Enthalpy
The ionisation enthalpy values as expected
from the general trends do not decrease
smoothly down the group. The decrease from
B to Al is associated with increase in size. The
observed discontinuity in the ionisation
enthalpy values between Al and Ga, and
between In and Tl are due to inability of d- and
f-electrons ,which have low screening effect, to
compensate the increase in nuclear charge.
The order of ionisation enthalpies, as
expected, is
i
H
1
<
i
H
2
<
i
H
3
. The sum of the
first three ionisation enthalpies for each of the
2019-20
318 CHEMISTRY
elements is very high. Effect of this will be
apparent when you study their chemical
properties.
11.1.4 Electronegativity
Down the group, electronegativity first
decreases from B to Al and then increases
marginally (Table 11.2). This is because of the
discrepancies in atomic size of the elements.
11.1.5 Physical Properties
Boron is non-metallic in nature. It is extremely
hard and black coloured solid. It exists in many
allotropic forms. Due to very strong crystalline
lattice, boron has unusually high melting point.
Rest of the members are soft metals with low
melting point and high electrical conductivity.
It is worthwhile to note that gallium with
unusually low melting point (303 K), could
exist in liquid state during summer. Its high
boiling point (2676 K) makes it a useful
material for measuring high temperatures.
Density of the elements increases down the
group from boron to thallium.
11.1.6 Chemical Properties
Oxidation state and trends in chemical
reactivity
Due to small size of boron, the sum of its first
three ionization enthalpies is very high. This
prevents it to form +3 ions and forces it to form
only covalent compounds. But as we move from
B to Al, the sum of the first three ionisation
enthalpies of Al considerably decreases, and
is therefore able to form Al
3+
ions. In fact,
aluminium is a highly electropositive metal.
However, down the group, due to poor
shielding effect of intervening d and f orbitals,
the increased effective nuclear charge holds ns
electrons tightly (responsible for inert pair
effect) and thereby, restricting their
participation in bonding. As a result of this,
only p-orbital electron may be involved in
bonding. In fact in Ga, In and Tl, both +1 and
+3 oxidation states are observed. The relative
stability of +1 oxidation state progressively
increases for heavier elements: Al<Ga<In<Tl. In
thallium +1 oxidation state is predominant
Table 11.2 Atomic and Physical Properties of Group 13 Elements
a
Metallic radius,
b
6-coordination,
c
Pauling scale,
Atomic number 5 13 31 49 81
Atomic mass
(g mol
–1
)
10.81 26.98 69.72 114.82 204.38
Electronic [He]2s
2
2p
1
[Ne]3s
2
3p
1
[Ar]3d
10
4s
2
4p
1
[Kr]4d
10
5s
2
5p
1
[Xe]4f
14
5d
10
6s
2
6p
1
Configuration
Atomic radius/pm
a
(88) 143 135 167 170
Ionic radius (27) 53.5 62.0 80.0 88.5
M
3+
/pm
b
Ionic radius --120 140 150
M
+
/pm
Ionization
i
H
1
801 577 579 558 589
enthalpy
i
H
2
2427 1816 1979 1820 1971
(kJ mol
–1
)
i
H
3
3659 2744 2962 2704 2877
Electronegativity
c
2.0 1.5 1.6 1.7 1.8
Density /g cm
–3
2.35 2.70 5.90 7.31 11.85
at 298 K
Melting point / K 2453 933 303 430 576
Boiling point / K 3923 2740 2676 2353 1730
E
/ V for (M
3+
/M) - –1.66 –0.56 –0.34 +1.26
E
/ V for (M
+
/M) - +0.55 -0.79(acid) –0.18 –0.34
–1.39(alkali)
Property
Element
Boron Aluminium Gallium Indium Thallium
B Al Ga In Tl
2019-20
319THE p-BLOCK ELEMENTS
whereas the +3 oxidation state is highly
oxidising in character. The compounds in
+1 oxidation state, as expected from energy
considerations, are more ionic than those in
+3 oxidation state.
In trivalent state, the number of electrons
around the central atom in a molecule
of the compounds of these elements
(e.g., boron in BF
3
) will be only six. Such
electron deficient molecules have tendency
to accept a pair of electrons to achieve stable
electronic configuration and thus, behave as
Lewis acids. The tendency to behave as Lewis
acid decreases with the increase in the size
down the group. BCl
3
easily accepts a lone pair
of electrons from ammonia to form BCl
3
NH
3
.
Solution
Standard electrode potential values for two
half cell reactions suggest that aluminium
has high tendency to make Al
3+
(aq) ions,
whereas Tl
3+
is not only unstable in
solution but is a powerful oxidising agent
also. Thus Tl
+
is more stable in solution
than Tl
3+
. Aluminium being able to form
+3 ions easily, is more electropositive than
thallium.
(i) Reactivity towards air
Boron is unreactive in crystalline form.
Aluminium forms a very thin oxide layer on
the surface which protects the metal from
further attack. Amorphous boron and
aluminium metal on heating in air form B
2
O
3
and Al
2
O
3
respectively. With dinitrogen at high
temperature they form nitrides.
(
)
(
)
(
)
() () ()
2 23
2
2E s 3O g 2E O s
2E s N g 2EN s
+ →
+ →
(E = element)
The nature of these oxides varies down the
group. Boron trioxide is acidic and reacts with
basic (metallic) oxides forming metal borates.
Aluminium and gallium oxides are amphoteric
and those of indium and thallium are basic in
their properties.
(ii) Reactivity towards acids and alkalies
Boron does not react with acids and alkalies
even at moderate temperature; but aluminium
dissolves in mineral acids and aqueous alkalies
and thus shows amphoteric character.
Aluminium dissolves in dilute HCl and
liberates dihydrogen.
2Al(s) + 6HCl (aq) 2Al
3+
(aq) + 6Cl
(aq)
+ 3H
2
(g)
However, concentrated nitric acid renders
aluminium passive by forming a protective
oxide layer on the surface.
Aluminium also reacts with aqueous alkali
and liberates dihydrogen.
2Al (s) + 2NaOH(aq) + 6H
2
O(l)
2 Na
+
[Al(OH)
4
]
(aq) + 3H
2
(g)
Sodium
tetrahydroxoaluminate(III)
In trivalent state most of the compounds
being covalent are hydrolysed in water. For
example, the trichlorides on hyrolysis in water
form tetrahedral
()
4
M OH


species; the
hybridisation state of element M is sp
3
.
Aluminium chloride in acidified aqueous
solution forms octahedral
()
3
2
6
Al H O
+


ion.
In this complex ion, the 3d orbitals of Al are
involved and the hybridisation state of Al is
sp
3
d
2
.
Problem 11.1
Standard electrode potential values, E
for Al
3+
/Al is –1.66 V and that of Tl
3+
/Tl
is +1.26 V. Predict about the formation of
M
3+
ion in solution and compare the
electropositive character of the two
metals.
AlCl
3
achieves stability by forming a dimer
2019-20
320 CHEMISTRY
(iii) Reactivity towards halogens
These elements react with halogens to form
trihalides (except TlI
3
).
2E(s) + 3 X
2
(g) 2EX
3
(s) (X = F, Cl, Br, I)
Problem 11.2
White fumes appear around the bottle of
anhydrous aluminium chloride. Give
reason.
Solution
Anhydrous aluminium chloride is
partially hydrolysed with atmospheric
moisture to liberate HCl gas. Moist HCl
appears white in colour.
11.2 IMPORTANT TRENDS AND
ANOMALOUS PROPERTIES OF
BORON
Certain important trends can be observed
in the chemical behaviour of group
13 elements. The tri-chlorides, bromides
and iodides of all these elements being
covalent in nature are hydrolysed in water.
Species like tetrahedral [M(OH)
4
]
and
octahedral [M(H
2
O)
6
]
3+
, except in boron, exist
in aqueous medium.
The monomeric trihalides, being electron
deficient, are strong Lewis acids. Boron
trifluoride easily reacts with Lewis bases such
as NH
3
to complete octet around boron.
3 3 3 3
FB :NH FB NH
It is due to the absence of d orbitals that
the maximum covalence of B is 4. Since the
d orbitals are available with Al and other
elements, the maximum covalence can be
expected beyond 4. Most of the other metal
halides (e.g., AlCl
3
) are dimerised through
halogen bridging (e.g., Al
2
Cl
6
). The metal
species completes its octet by accepting
electrons from halogen in these halogen
bridged molecules.
Problem 11.3
Boron is unable to form BF
6
3–
ion. Explain.
Solution
Due to non-availability of d orbitals, boron
is unable to expand its octet. Therefore,
the maximum covalence of boron cannot
exceed 4.
11.3 SOME IMPORTANT COMPOUNDS OF
BORON
Some useful compounds of boron are borax,
orthoboric acid and diborane. We will briefly
study their chemistry.
11.3.1 Borax
It is the most important compound of boron.
It is a white crystalline solid of formula
Na
2
B
4
O
7
⋅⋅
10H
2
O. In fact it contains the
tetranuclear units
()
2
45
4
B O OH


and correct
formula; therefore, is Na
2
[B
4
O
5
(OH)
4
].8H
2
O.
Borax dissolves in water to give an alkaline
solution.
Na
2
B
4
O
7
+ 7H
2
O 2NaOH + 4H
3
BO
3
Orthoboric acid
On heating, borax first loses water
molecules and swells up. On further heating it
turns into a transparent liquid, which solidifies
into glass like material known as borax
bead.
Na
2
B
4
O
7
.10H
2
O
→
Na
2
B
4
O
7
→
2NaBO
2
Sodium + B
2
O
3
metaborate Boric
anhydride
The metaborates of many transition metals
have characteristic colours and, therefore,
borax bead test can be used to identify them
in the laboratory. For example, when borax is
heated in a Bunsen burner flame with CoO on
a loop of platinum wire, a blue coloured
Co(BO
2
)
2
bead is formed.
11.3.2 Orthoboric acid
Orthoboric acid, H
3
BO
3
is a white crystalline
solid, with soapy touch. It is sparingly soluble
in water but highly soluble in hot water. It can
be prepared by acidifying an aqueous solution
of borax.
Na
2
B
4
O
7
+ 2HCl + 5H
2
O