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, C≡C,
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
N≡N) and to other second row elements (e.g.,
C=O, C=N, C≡N, 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
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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 → 2NaCl + 4B(OH)
3
It is also formed by the hydrolysis (reaction
with water or dilute acid) of most boron
compounds (halides, hydrides, etc.). It has a
2019-20
321THE p-BLOCK ELEMENTS
Problem 11.4
Why is boric acid considered as a weak
acid?
Solution
Because it is not able to release H
+
ions
on its own. It receives OH
–
ions from water
molecule to complete its octet and in turn
releases H
+
ions.
11.3.3 Diborane, B
2
H
6
The simplest boron hydride known, is
diborane. It is prepared by treating boron
trifluoride with LiAlH
4
in diethyl ether.
4BF
3
+ 3 LiAlH
4
→ 2B
2
H
6
+ 3LiF + 3AlF
3
A convenient laboratory method for the
preparation of diborane involves the oxidation
of sodium borohydride with iodine.
2NaBH
4
+ I
2
→ B
2
H
6
+ 2NaI + H
2
Diborane is produced on an industrial scale
by the reaction of BF
3
with sodium hydride.
3 26
450K
2BF 6NaH B H 6NaF
+ → +
Diborane is a colourless, highly toxic gas
with a b.p. of 180 K. Diborane catches fire
spontaneously upon exposure to air. It burns
in oxygen releasing an enormous amount of
energy.
26 2 23 2
1
B H +3O B O + 3H O;
1976 kJ mol
−
→
∆ =−
c
H
Most of the higher boranes are also
spontaneously flammable in air. Boranes are
readily hydrolysed by water to give boric acid.
B
2
H
6
(g) + 6H
2
O(l) → 2B(OH)
3
(aq) + 6H
2
(g)
Diborane undergoes cleavage reactions
with Lewis bases(L) to give borane adducts,
BH
3
⋅⋅
⋅⋅
⋅L
B
2
H
6
+ 2 NMe
3
→ 2BH
3
⋅⋅
⋅⋅
⋅NMe
3
B
2
H
6
+ 2 CO → 2BH
3
⋅⋅
⋅⋅
⋅CO
Reaction of ammonia with diborane gives
initially B
2
H
6
.2NH
3
which is formulated as
[BH
2
(NH
3
)
2
]
+
[BH
4
]
–
; further heating gives
borazine, B
3
N
3
H
6
known as “inorganic
benzene” in view of its ring structure with
alternate BH and NH groups.
–
26 3 2 32 4
336 2
+
Heat
3B H +6NH 3[BH (NH ) ] [BH ]
2B N H +12H
→
→
The structure of diborane is shown in
Fig.11.2(a). The four terminal hydrogen atoms
and the two boron atoms lie in one plane.
Above and below this plane, there are two
bridging hydrogen atoms. The four terminal
B-H bonds are regular two centre-two electron
bonds while the two bridge (B-H-B) bonds are
different and can be described in terms of three
Fig.11.2(a) The structure of diborane, B
2
H
6
Fig. 11. 1 Structure of boric acid; the dotted lines
represent hydrogen bonds
layer structure in which planar BO
3
units are
joined by hydrogen bonds as shown in
Fig. 11.1.
Boric acid is a weak monobasic acid. It is
not a protonic acid but acts as a Lewis acid
by accepting electrons from a hydroxyl ion:
B(OH)
3
+ 2HOH → [B(OH)
4
]
–
+ H
3
O
+
On heating, orthoboric acid above 370K
forms metaboric acid, HBO
2
which on further
heating yields boric oxide, B
2
O
3
.
H
3
BO
3
∆
→
HBO
2
∆
→
B
2
O
3
2019-20
322 CHEMISTRY
centre–two electron bonds shown in
Fig.11.2 (b).
Boron also forms a series of hydridoborates;
the most important one is the tetrahedral [BH
4
]
–
ion. Tetrahydridoborates of several metals are
known. Lithium and sodium tetra-
hydridoborates, also known as borohydrides,
are prepared by the reaction of metal hydrides
with B
2
H
6
in diethyl ether.
2MH + B
2
H
6
→ 2 M
+
[BH
4
]
–
(M = Li or Na)
orthoboric acid is generally used as a mild
antiseptic.
Aluminium is a bright silvery-white metal,
with high tensile strength. It has a high
electrical and thermal conductivity. On a
weight-to-weight basis, the electrical
conductivity of aluminium is twice that of
copper. Aluminium is used extensively in
industry and everyday life. It forms alloys with
Cu, Mn, Mg, Si and Zn. Aluminium and its
alloys can be given shapes of pipe, tubes,
rods, wires, plates or foils and, therefore, find
uses in packing, utensil making,
construction, aeroplane and transportation
industry. The use of aluminium and its
compounds for domestic purposes is now
reduced considerably because of their toxic
nature.
11.5 GROUP 14 ELEMENTS: THE CARBON
FAMILY
Carbon, silicon, germanium, tin lead and
flerovium are the members of group 14. Carbon
is the seventeenth most abundant element by
mass in the earth’s crust. It is widely
distributed in nature in free as well as in the
combined state. In elemental state it is available
as coal, graphite and diamond; however, in
combined state it is present as metal
carbonates, hydrocarbons and carbon dioxide
gas (0.03%) in air. One can emphatically say
that carbon is the most versatile element in the
world. Its combination with other elements
such as dihydrogen, dioxygen, chlorine and
sulphur provides an astonishing array of
materials ranging from living tissues to drugs
and plastics. Organic chemistry is devoted to
carbon containing compounds. It is an
essential constituent of all living organisms.
Naturally occurring carbon contains two stable
isotopes:
12
C and
13
C. In addition to these, third
isotope,
14
C is also present. It is a radioactive
isotope with half-life 5770 years and used for
radiocarbon dating. Silicon is the second
(27.7 % by mass) most abundant element on
the earth’s crust and is present in nature in
the form of silica and silicates. Silicon is a very
important component of ceramics, glass and
cement. Germanium exists only in traces. Tin
Both LiBH
4
and NaBH
4
are used as
reducing agents in organic synthesis. They are
useful starting materials for preparing other
metal borohydrides.
11.4 USES OF BORON AND ALUMINIUM
AND THEIR COMPOUNDS
Boron being extremely hard refractory solid of
high melting point, low density and very low
electrical conductivity, finds many
applications. Boron fibres are used in making
bullet-proof vest and light composite material
for aircraft. The boron-10 (
10
B) isotope has high
ability to absorb neutrons and, therefore,
metal borides are used in nuclear industry as
protective shields and control rods. The main
industrial application of borax and boric acid
is in the manufacture of heat resistant glasses
(e.g., Pyrex), glass-wool and fibreglass. Borax
is also used as a flux for soldering metals, for
heat, scratch and stain resistant glazed coating
to earthenwares and as constituent of
medicinal soaps. An aqueous solution of
Fig.11.2(b) Bonding in diborane. Each B atom
uses sp
3
hybrids for bonding. Out
of the four sp
3
hybrids on each B
atom, one is without an electron
shown in broken lines. The terminal
B-H bonds are normal 2-centre-2-
electron bonds but the two bridge
bonds are 3-centre-2-electron bonds.
The 3-centre-2-electron bridge bonds
are also referred to as banana bonds.
2019-20
323THE p-BLOCK ELEMENTS
occurs mainly as cassiterite, SnO
2
and lead as
galena, PbS. Flerovium is synthetically
prepared radioactive element
Ultrapure form of germanium and silicon
are used to make transistors and
semiconductor devices.
Symbol of Flerovium is Fl. It has atomic
number 114, atomic mass 289 gmol
-1
and
electronic configuration [Rn] 5f
14
6d
10
7s
2
7p
2
.
It has been prepared only in small amount.
Its half life is short and its chemistry has not
been established yet. The important atomic
and physical properties along with their
electronic configuration of the elements of
group 14 leaving flerovium are given in
Table 11.3. Some of the atomic, physical and
chemical properties are discussed below:
11.5.1 Electronic Configuration
The valence shell electronic configuration of
these elements is ns
2
np
2
. The inner core of the
electronic configuration of elements in this
group also differs.
11.5.2 Covalent Radius
There is a considerable increase in covalent
radius from C to Si, thereafter from Si to Pb a
small increase in radius is observed. This is
due to the presence of completely filled d and f
orbitals in heavier members.
11.5.3 Ionization Enthalpy
The first ionization enthalpy of group 14
members is higher than the corresponding
members of group 13. The influence of inner
core electrons is visible here also. In general the
ionisation enthalpy decreases down the group.
Small decrease in ∆
i
H from Si to Ge to Sn and
slight increase in ∆
i
H from Sn to Pb is the
consequence of poor shielding effect of
intervening d and f orbitals and increase in size
of the atom.
11.5.4 Electronegativity
Due to small size, the elements of this group
are slightly more electronegative than group
13 elements. The electronegativity values for
elements from Si to Pb are almost the same.
Table 11.3 Atomic and Physical Properties of Group 14 Elements
a
for M
IV
oxidation state;
b
6–coordination;
c
Pauling scale;
d
293 K;
e
for diamond; for graphite, density is
2.22;
f
β
-form (stable at room temperature)
Atomic Number 6 14 32 50 82
Atomic mass (g mol
–1
) 12.01 28.09 72.60 118.71 207.2
Electronic [He]2s
2
2p
2
[Ne]3s
2
3p
2
[Ar]3d
10
4s
2
4p
2
[Kr]4d
10
5s
2
5p
2
[Xe]4f
14
5d
10
6s
2
6p
2
configuration
Covalent radius/pm
a
77 118 122 140 146
Ionic radius M
4+
/pm
b
– 40 53 69 78
Ionic radius M
2+
/pm
b
– – 73 118 119
Ionization ∆
i
H
1
1086 786 761 708 715
enthalpy/ ∆
i
H
2
2352 1577 1537 1411 1450
kJ mol
–1
∆
i
H
3
4620 3228 3300 2942 3081
∆
i
H
4
6220 4354 4409 3929 4082
Electronegativity
c
2.5 1.8 1.8 1.8 1.9
Density
d
/g cm
–3
3.51
e
2.34 5.32 7.26
f
11.34
Melting point/K 4373 1693 1218 505 600
Boiling point/K – 3550 3123 2896 2024
Electrical resistivity/ 10
14
–10
16
50 50 10
–5
2 × 10
–5
ohm cm (293 K)
Carbon Silicon Germanium Tin Lead
C Si Ge Sn Pb
Element
Property
2019-20
324 CHEMISTRY
11.5.5 Physical Properties
All members of group14 are solids. Carbon and
silicon are non-metals, germanium is a metalloid,
whereas tin and lead are soft metals with low
melting points. Melting points and boiling points
of group 14 elements are much higher than those
of corresponding elements of group 13.
11.5.6 Chemical Properties
Oxidation states and trends in chemical
reactivity
The group 14 elements have four electrons in
outermost shell. The common oxidation states
exhibited by these elements are +4 and +2.
Carbon also exhibits negative oxidation states.
Since the sum of the first four ionization
enthalpies is very high, compounds in +4
oxidation state are generally covalent in nature.
In heavier members the tendency to show +2
oxidation state increases in the sequence
Ge<Sn<Pb. It is due to the inability of ns
2
electrons of valence shell to participate in
bonding. The relative stabilities of these two
oxidation states vary down the group. Carbon
and silicon mostly show +4 oxidation state.
Germanium forms stable compounds in +4
state and only few compounds in +2 state. Tin
forms compounds in both oxidation states (Sn
in +2 state is a reducing agent). Lead
compounds in +2 state are stable and in +4
state are strong oxidising agents. In tetravalent
state the number of electrons around the
central atom in a molecule (e.g., carbon in CCl
4
)
is eight. Being electron precise molecules, they
are normally not expected to act as electron
acceptor or electron donor species. Although
carbon cannot exceed its covalence more than
4, other elements of the group can do so. It is
because of the presence of d orbital in them.
Due to this, their halides undergo hydrolysis
and have tendency to form complexes by
accepting electron pairs from donor species. For
example, the species like, SiF
6
2–
, [GeCl
6
]
2–
,
[Sn(OH)
6
]
2–
exist where the hybridisation of the
central atom is sp
3
d
2
.
(i) Reactivity towards oxygen
All members when heated in oxygen form
oxides. There are mainly two types of oxides,
i.e., monoxide and dioxide of formula MO and
MO
2
respectively. SiO only exists at high
temperature. Oxides in higher oxidation states
of elements are generally more acidic than
those in lower oxidation states. The dioxides
— CO
2
, SiO
2
and GeO
2
are acidic, whereas
SnO
2
and PbO
2
are amphoteric in nature.
Among monoxides, CO is neutral, GeO is
distinctly acidic whereas SnO and PbO are
amphoteric.
Problem 11.5
Select the member(s) of group 14 that
(i) forms the most acidic dioxide, (ii) is
commonly found in +2 oxidation state,
(iii) used as semiconductor.
Solution
(i) carbon (ii) lead
(iii) silicon and germanium
(ii) Reactivity towards water
Carbon, silicon and germanium are not
affected by water. Tin decomposes steam to
form dioxide and dihydrogen gas.
2 22
Sn + 2H O SnO + 2H
∆
→
Lead is unaffected by water, probably
because of a protective oxide film formation.
(iii) Reactivity towards halogen
These elements can form halides of formula
MX
2
and MX
4
(where X = F, Cl, Br, I). Except
carbon, all other members react directly with
halogen under suitable condition to make
halides. Most of the MX
4
are covalent in nature.
The central metal atom in these halides
undergoes sp
3
hybridisation and the molecule
is tetrahedral in shape. Exceptions are SnF
4
and PbF
4
, which are ionic in nature. PbI
4
does
not exist because Pb—I bond initially formed
during the reaction does not release enough
energy to unpair 6s
2
electrons and excite one
of them to higher orbital to have four unpaired
electrons around lead atom. Heavier members
Ge to Pb are able to make halides of formula
MX
2
. Stability of dihalides increases down the
group. Considering the thermal and chemical
stability, GeX
4
is more stable than GeX
2
,
whereas PbX
2
is more than PbX
4
. Except CCl
4
,
other tetrachlorides are easily hydrolysed
2019-20
325THE p-BLOCK ELEMENTS
by water because the central atom can
accommodate the lone pair of electrons from
oxygen atom of water molecule in d orbital.
Hydrolysis can be understood by taking
the example of SiCl
4.
It
undergoes hydrolysis
by initially accepting lone pair of electrons
from water molecule in d orbitals of Si, finally
leading to the formation of Si(OH)
4
as shown
below :
Carbon also has unique ability to form
p
π
– p
π
multiple bonds with itself and with other
atoms of small size and high electronegativity.
Few examples of multiple bonding are: C=C,
C ≡ C, C = O, C = S, and C ≡ N. Heavier elements
do not form p
π
– p
π
bonds because their atomic
orbitals are too large and diffuse to have
effective overlapping.
Carbon atoms have the tendency to link
with one another through covalent bonds to
form chains and rings. This property is called
catenation. This is because C—C bonds are
very strong. Down the group the size increases
and electronegativity decreases, and, thereby,
tendency to show catenation decreases. This
can be clearly seen from bond enthalpies
values. The order of catenation is C > > Si >
Ge ≈ Sn. Lead does not show catenation.
Bond Bond enthalpy / kJ mol
–1
C—C 348
Si —Si 297
Ge—Ge 260
Sn—Sn 240
Due to property of catenation and p
π
– p
π
bond formation, carbon is able to show
allotropic forms.
11.7 ALLOTROPES OF CARBON
Carbon exhibits many allotropic forms; both
crystalline as well as amorphous. Diamond
and graphite are two well-known crystalline
forms of carbon. In 1985, third form of carbon
known as fullerenes was discovered by
H.W.Kroto, E.Smalley and R.F.Curl. For this
discovery they were awarded the Nobel Prize
in 1996.
11.7.1 Diamond
It has a crystalline lattice. In diamond each
carbon atom undergoes sp
3
hybridisation and
linked to four other carbon atoms by using
hybridised orbitals in tetrahedral fashion. The
C–C bond length is 154 pm. The structure
extends in space and produces a rigid three-
dimensional network of carbon atoms. In this
Problem 11. 6
[SiF
6
]
2–
is known whereas [SiCl
6
]
2–
not.
Give possible reasons.
Solution
The main reasons are :
(i) six large chloride ions cannot be
accommodated around Si
4+
due to
limitation of its size.
(ii) interaction between lone pair of
chloride ion and Si
4+
is not very strong.
11.6 IMPORTANT TRENDS AND
ANOMALOUS BEHAVIOUR OF
CARBON
Like first member of other groups, carbon
also differs from rest of the members of its
group. It is due to its smaller size, higher
electronegativity, higher ionisation enthalpy
and unavailability of d orbitals.
In carbon, only s and p orbitals are
available for bonding and, therefore, it can
accommodate only four pairs of electrons
around it. This would limit the maximum
covalence to four whereas other members can
expand their covalence due to the presence of
d orbitals.
2019-20
326 CHEMISTRY
structure (Fig. 11.3) directional covalent bonds
are present throughout the lattice.
It is very difficult to break extended covalent
bonding and, therefore, diamond is a hardest
substance on the earth. It is used as an
abrasive for sharpening hard tools, in making
dyes and in the manufacture of tungsten
filaments for electric light bulbs.
Problem 11.7
Diamond is covalent, yet it has high
melting point. Why ?
Solution
Diamond has a three-dimensional
network involving strong C—C bonds,
which are very difficult to break and, in
turn has high melting point.
11.7.2 Graphite
Graphite has layered structure (Fig.11.4).
Layers are held by van der Waals forces and
distance between two layers is 340 pm. Each
layer is composed of planar hexagonal rings
of carbon atoms. C—C bond length within the
layer is 141.5 pm. Each carbon atom in
hexagonal ring undergoes sp
2
hybridisation
and makes three sigma bonds with three
neighbouring carbon atoms. Fourth electron
forms a
π
bond. The electrons are delocalised
over the whole sheet. Electrons are mobile and,
therefore, graphite conducts electricity along
the sheet. Graphite cleaves easily between the
layers and, therefore, it is very soft and slippery.
For this reason graphite is used as a dry
lubricant in machines running at high
temperature, where oil cannot be used as a
lubricant.
11.7.3 Fullerenes
Fullerenes are made by the heating of graphite
in an electric arc in the presence of inert gases
such as helium or argon. The sooty material
formed by condensation of vapourised C
n
small
molecules consists of mainly C
60
with smaller
quantity of C
70
and traces of fullerenes
consisting of even number of carbon atoms up
to 350 or above. Fullerenes are the only pure
form of carbon because they have smooth
structure without having ‘dangling’ bonds.
Fullerenes are cage like molecules. C
60
molecule has a shape like soccer ball and
called Buckminsterfullerene (Fig. 11.5).
It contains twenty six- membered rings and
twelve five-membered rings. A six membered
ring is fused with six or five membered rings
but a five membered ring can only fuse with
six membered rings. All the carbon atoms are
equal and they undergo sp
2
hybridisation.
Each carbon atom forms three sigma bonds
with other three carbon atoms. The remaining
electron at each carbon is delocalised in
Fig. 11.3 The structure of diamond Fig 11.4 The structure of graphite
2019-20
327THE p-BLOCK ELEMENTS
molecular orbitals, which in turn give aromatic
character to molecule. This ball shaped
molecule has 60 vertices and each one is
occupied by one carbon atom and it also
contains both single and double bonds with
C–C distances of 143.5 pm and 138.3 pm
respectively. Spherical fullerenes are also called
bucky balls in short.
filters to remove organic contaminators and in
airconditioning system to control odour.
Carbon black is used as black pigment in
black ink and as filler in automobile tyres. Coke
is used as a fuel and largely as a reducing
agent in metallurgy. Diamond is a precious
stone and used in jewellery. It is measured in
carats (1 carat = 200 mg).
11.8 SOME IMPORTANT COMPOUNDS OF
CARBON AND SILICON
Oxides of Carbon
Two important oxides of carbon are carbon
monoxide, CO and carbon dioxide, CO
2
.
11.8.1 Carbon Monoxide
Direct oxidation of C in limited supply of
oxygen or air yields carbon monoxide.
2
2C(s) O (g) 2CO(g)
∆
+ →
On small scale pure CO is prepared by
dehydration of formic acid with concentrated
H
2
SO
4
at 373 K
2
4
2
373K
conc.H SO
HCOOH H O + CO
→
On commercial scale it is prepared by the
passage of steam over hot coke. The mixture
of CO and H
2
thus produced is known as water
gas or synthesis gas.
() () () ()
2 2
473 1273K
Cs HOg CO g H g
Water gas
−
+ → +
When air is used instead of steam, a mixture
of CO and N
2
is produced, which is called
producer gas.
2 2
2
1273K
2C(s) O (g) 4N (g) 2CO(g)
4N (g)
+ + →
+
Producer gas
Water gas and producer gas are very
important industrial fuels. Carbon monoxide
in water gas or producer gas can undergo
further combustion forming carbon dioxide
with the liberation of heat.
Carbon monoxide is a colourless,
odourless and almost water insoluble gas. It
is a powerful reducing agent and reduces
almost all metal oxides other than those of the
alkali and alkaline earth metals, aluminium
and a few transition metals. This property of
Fig.11.5 The structure of C
60
, Buckminster-
fullerene : Note that molecule has the
shape of a soccer ball (football).
It is very important to know that graphite
is thermodynamically most stable allotrope of
carbon and, therefore, ∆
f
H
of graphite is taken
as zero. ∆
f
H
values of diamond and fullerene,
C
60
are 1.90 and 38.1 kJ mol
–1
, respectively.
Other forms of elemental carbon like carbon
black, coke, and charcoal are all impure forms
of graphite or fullerenes. Carbon black is
obtained by burning hydrocarbons in a limited
supply of air. Charcoal and coke are obtained
by heating wood or coal respectively at high
temperatures in the absence of air.
11.7.4 Uses of Carbon
Graphite fibres embedded in plastic material
form high strength, lightweight composites.
The composites are used in products such as
tennis rackets, fishing rods, aircrafts and
canoes. Being good conductor, graphite is used
for electrodes in batteries and industrial
electrolysis. Crucibles made from graphite are
inert to dilute acids and alkalies. Being highly
porous, activated charcoal is used in
adsorbing poisonous gases; also used in water
2019-20
328 CHEMISTRY
CO is used in the extraction of many metals
from their oxides ores.
() () () ()
() () () ()
23 2
2
Fe O s 3CO g 2Fe s 3CO g
ZnO s CO g Zn s CO g
∆
∆
+ → +
+ → +
In CO molecule, there are one sigma and
two
π
bonds between carbon and oxygen,
:C ≡ O: . Because of the presence of a lone pair
on carbon, CO molecule acts as a donor and
reacts with certain metals when heated to form
metal carbonyls. The highly poisonous
nature of CO arises because of its ability to
form a complex with haemoglobin, which
is about 300 times more stable than the
oxygen-haemoglobin complex. This prevents
haemoglobin in the red blood corpuscles from
carrying oxygen round the body and ultimately
resulting in death.
11.8.2 Carbon Dioxide
It is prepared by complete combustion of
carbon and carbon containing fuels in excess
of air.
2 2
C(s) O (g) CO (g)
∆
+ →
4 2 2 2
CH (g) 2O (g) CO (g) 2H O(g)
∆
+ → +
In the laboratory it is conveniently
prepared by the action of dilute HCl on calcium
carbonate.
CaCO
3
(s) + 2HCl (aq) → CaCl
2
(aq) + CO
2
(g) +
H
2
O(l)
On commercial scale it is obtained by
heating limestone.
It is a colourless and odourless gas. Its low
solubility in water makes it of immense bio-
chemical and geo-chemical importance. With
water, it forms carbonic acid, H
2
CO
3
which is
a weak dibasic acid and dissociates in two
steps:
H
2
CO
3
(aq) + H
2
O(l) HCO
3
–
(aq) + H
3
O
+
(aq)
HCO
3
–
(aq) + H
2
O(l) CO
3
2–
(aq) + H
3
O
+
(aq)
H
2
CO
3
/HCO
3
–
buffer system helps to
maintain pH of blood between 7.26 to 7.42.
Being acidic in nature, it combines with alkalies
to form metal carbonates.
Carbon dioxide, which is normally present
to the extent of ~ 0.03 % by volume in the
atmosphere, is removed from it by the process
known as photosynthesis. It is the process
by which green plants convert atmospheric
CO
2
into carbohydrates such as glucose. The
overall chemical change can be expressed as:
→
22 6 12 6 2
2
h
Chlorophyll
6CO +12H O C H O + 6O
+ 6H O
ν
By this process plants make food for
themselves as well as for animals and human
beings. Unlike CO, it is not poisonous. But the
increase in combustion of fossil fuels and
decomposition of limestone for cement
manufacture in recent years seem to increase
the CO
2
content of the atmosphere. This may
lead to increase in green house effect and
thus, raise the temperature of the atmosphere
which might have serious consequences.
Carbon dioxide can be obtained as a solid
in the form of dry ice by allowing the liquified
CO
2
to expand rapidly. Dry ice is used as a
refrigerant for ice-cream and frozen food.
Gaseous CO
2
is extensively used to carbonate
soft drinks. Being heavy and non-supporter
of combustion it is used as fire extinguisher. A
substantial amount of CO
2
is used to
manufacture urea.
In CO
2
molecule carbon atom undergoes
sp hybridisation. Two sp hybridised orbitals
of carbon atom overlap with two p orbitals of
oxygen atoms to make two sigma bonds while
other two electrons of carbon atom are involved
in p
π
– p
π
bonding with oxygen atom. This
results in its linear shape [with both C–O bonds
of equal length (115 pm)] with no dipole
moment. The resonance structures are shown
below:
Resonance structures of carbon dioxide
11.8.3 Silicon Dioxide, SiO
2
95% of the earth’s crust is made up of silica
and silicates. Silicon dioxide, commonly known
as silica, occurs in several crystallographic
forms. Quartz, cristobalite and tridymite are
some of the crystalline forms of silica, and they
are interconvertable at suitable temperature.
Silicon dioxide is a covalent, three-dimensional
2019-20
329THE p-BLOCK ELEMENTS
network solid in which each silicon atom is
covalently bonded in a tetrahedral manner to
four oxygen atoms. Each oxygen atom in turn
covalently bonded to another silicon atoms as
shown in diagram (Fig 11.6 ). Each corner is
shared with another tetrahedron. The entire
crystal may be considered as giant molecule
in which eight membered rings are formed with
alternate silicon and oxygen atoms.
substituted chlorosilane of formula MeSiCl
3
,
Me
2
SiCl
2
, Me
3
SiCl with small amount of Me
4
Si
are formed. Hydrolysis of dimethyl-
dichlorosilane, (CH
3
)
2
SiCl
2
followed by
condensation polymerisation yields straight
chain polymers.
The chain length of the polymer can be
controlled by adding (CH
3
)
3
SiCl which blocks
the ends as shown below :
Fig. 11.6 Three dimensional structure of SiO
2
Silica in its normal form is almost non-
reactive because of very high Si— O bond
enthalpy. It resists the attack by halogens,
dihydrogen and most of the acids and metals
even at elevated temperatures. However, it is
attacked by HF and NaOH.
SiO
2
+ 2NaOH → Na
2
SiO
3
+ H
2
O
SiO
2
+ 4HF → SiF
4
+ 2H
2
O
Quartz is extensively used as a piezoelectric
material; it has made possible to develop extremely
accurate clocks, modern radio and television
broadcasting and mobile radio communications.
Silica gel is used as a drying agent and as a support
for chromatographic materials and catalysts.
Kieselghur, an amorphous form of silica is used
in filtration plants.
11.8.4 Silicones
They are a group of organosilicon polymers,
which have (
R
2
SiO) as a repeating unit. The
starting materials for the manufacture of
silicones are alkyl or aryl substituted silicon
chlorides, R
n
SiCl
(4–n)
, where R is alkyl or aryl
group. When methyl chloride reacts with
silicon in the presence of copper as a catalyst
at a temperature 573K various types of methyl
2019-20
330 CHEMISTRY
Silicones being surrounded by non-polar
alkyl groups are water repelling in nature.
They have in general high thermal stability,
high dielectric strength and resistance to
oxidation and chemicals. They have wide
applications. They are used as sealant, greases,
electrical insulators and for water proofing of
fabrics. Being biocompatible they are also used
in surgical and cosmetic plants.
Problem: 11.8
What are silicones ?
Solution
Simple silicones consist of
chains in which alkyl or phenyl groups
occupy the remaining bonding positions
on each silicon. They are hydrophobic
(water repellant) in nature.
11.8.5 Silicates
A large number of silicates minerals exist in
nature. Some of the examples are feldspar,
zeolites, mica and asbestos. The basic
structural unit of silicates is SiO
4
4–
(Fig.11.7)
in which silicon atom is bonded to four
oxygen atoms in tetrahedron fashion. In
silicates either the discrete unit is present or
a number of such units are joined together
via corners by sharing 1,2,3 or 4 oxygen
atoms per silicate units. When silicate units
are linked together, they form chain, ring,
sheet or three-dimensional structures.
Negative charge on silicate structure is
neutralised by positively charged metal ions.
If all the four corners are shared with other
tetrahedral units, three-dimensional network
is formed.
Two important man-made silicates are
glass and cement.
11.8.6 Zeolites
If aluminium atoms replace few silicon atoms
in three-dimensional network of silicon dioxide,
overall structure known as aluminosilicate,
acquires a negative charge. Cations such as
Na
+
, K
+
or Ca
2+
balance the negative charge.
Examples are feldspar and zeolites. Zeolites are
widely used as a catalyst in petrochemical
industries for cracking of hydrocarbons and
isomerisation, e.g., ZSM-5 (A type of zeolite)
used to convert alcohols directly into gasoline.
Hydrated zeolites are used as ion exchangers
in softening of “hard” water.
SUMMARY
p-Block of the periodic table is unique in terms of having all types of elements – metals,
non-metals and metalloids. There are six groups of p-block elements in the periodic
table numbering from 13 to 18. Their valence shell electronic configuration is ns
2
np
1–6
(except for He). Differences in the inner core of their electronic configuration greatly
influence their physical and chemical properties. As a consequence of this, a lot of
variation in properties among these elements is observed. In addition to the group oxidation
state, these elements show other oxidation states differing from the total number of valence
electrons by unit of two. While the group oxidation state is the most stable for the lighter
elements of the group, lower oxidation states become progressively more stable for the
heavier elements. The combined effect of size and availability of d orbitals considerably
(a) (b)
Fig. 11.7 (a) Tetrahedral structure of SiO
4
4–
anion; (b) Representation of SiO
4
4–
unit
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331THE p-BLOCK ELEMENTS
influences the ability of these elements to form
π
-bonds. While the lighter elements form
p
ππ
ππ
π
–p
ππ
ππ
π
bonds, the heavier ones form d
ππ
ππ
π
–p
ππ
ππ
π
or d
ππ
ππ
π
–d
ππ
ππ
π
bonds. Absence of d orbital in
second period elements limits their maximum covalence to 4 while heavier ones can
exceed this limit.
Boron is a typical non-metal and the other members are metals. The availability of 3
valence electrons (2s
2
2p
1
) for covalent bond formation using four orbitals (2s, 2p
x,
2p
y
and
2p
z
) leads to the so called electron deficiency in boron compounds. This deficiency
makes them good electron acceptor and thus boron compounds behave as Lewis acids.
Boron forms covalent molecular compounds with dihydrogen as boranes, the simplest of
which is diborane, B
2
H
6
. Diborane contains two bridging hydrogen atoms between two
boron atoms; these bridge bonds are considered to be three-centre two-electron bonds.
The important compounds of boron with dioxygen are boric acid and borax. Boric acid,
B(OH)
3
is a weak monobasic acid; it acts as a Lewis acid by accepting electrons from
hydroxyl ion. Borax is a white crystalline solid of formula Na
2
[B
4
O
5
(OH)
4
]·8H
2
O. The borax
bead test gives characteristic colours of transition metals.
Aluminium exhibits +3 oxidation state. With heavier elements +1 oxidation state gets
progressively stabilised on going down the group. This is a consequence of the so called
inert pair effect.
Carbon is a typical non-metal forming covalent bonds employing all its four valence
electrons (2s
2
2p
2
). It shows the property of catenation, the ability to form chains or
rings, not only with C–C single bonds but also with multiple bonds (C=C or C≡C). The
tendency to catenation decreases as C>>Si>Ge
~ Sn > Pb. Carbon provides one of the
best examples of allotropy. Three important allotropes of carbon are diamond, graphite
and fullerenes. The members of the carbon family mainly exhibit +4 and +2 oxidation
states; compouds in +4 oxidation states are generally covalent in nature. The tendency
to show +2 oxidation state increases among heavier elements. Lead in +2 state is stable
whereas in +4 oxidation state it is a strong oxidising agent. Carbon also exhibits negative
oxidation states. It forms two important oxides: CO and CO
2
. Carbon monoxide is neutral
whereas CO
2
is acidic in nature. Carbon monoxide having lone pair of electrons on C
forms metal carbonyls. It is deadly poisonous due to higher stability of its haemoglobin
complex as compared to that of oxyhaemoglobin complex. Carbon dioxide as such is not
toxic. However, increased content of CO
2
in atmosphere due to combustion of fossil fuels
and decomposition of limestone is feared to cause increase in ‘green house effect’. This,
in turn, raises the temperature of the atmosphere and causes serious complications.
Silica, silicates and silicones are important class of compounds and find applications
in industry and technology.
EXERCISES
11.1 Discuss the pattern of variation in the oxidation states of
(i) B to Tl and (ii) C to Pb.
11.2 How can you explain higher stability of BCl
3
as compared to TlCl
3
?
11.3 Why does boron triflouride behave as a Lewis acid ?
11.4 Consider the compounds, BCl
3
and CCl
4
. How will they behave with
water ? Justify.
11.5 Is boric acid a protic acid ? Explain.
11.6 Explain what happens when boric acid is heated .
11.7 Describe the shapes of BF
3
and BH
4
–
. Assign the hybridisation of boron in
these species.
11.8 Write reactions to justify amphoteric nature of aluminium.
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332 CHEMISTRY
11.9 What are electron deficient compounds ? Are BCl
3
and SiCl
4
electron
deficient species ? Explain.
11.10 Write the resonance structures of CO
3
2–
and HCO
3
–
.
11.11 What is the state of hybridisation of carbon in (a) CO
3
2–
(b) diamond
(c) graphite?
11.12 Explain the difference in properties of diamond and graphite on the basis
of their structures.
11.13 Rationalise the given statements and give chemical reactions :
• Lead(II) chloride reacts with Cl
2
to give PbCl
4
.
• Lead(IV) chloride is highly unstable towards heat.
• Lead is known not to form an iodide, PbI
4
.
11.14 Suggest reasons why the B–F bond lengths in BF
3
(130 pm) and BF
4
–
(143 pm) differ.
11.15 If B–Cl bond has a dipole moment, explain why BCl
3
molecule has zero
dipole moment.
11.16 Aluminium trifluoride is insoluble in anhydrous HF but dissolves on
addition of NaF. Aluminium trifluoride precipitates out of the resulting
solution when gaseous BF
3
is bubbled through. Give reasons.
11.17 Suggest a reason as to why CO is poisonous.
11.18 How is excessive content of CO
2
responsible for global warming ?
11.19 Explain structures of diborane and boric acid.
11.20 What happens when
(a) Borax is heated strongly,
(b) Boric acid is added to water,
(c) Aluminium is treated with dilute NaOH,
(d) BF
3
is reacted with ammonia ?
11.21 Explain the following reactions
(a) Silicon is heated with methyl chloride at high temperature in the
presence of copper;
(b) Silicon dioxide is treated with hydrogen fluoride;
(c) CO is heated with ZnO;
(d) Hydrated alumina is treated with aqueous NaOH solution.
11.22 Give reasons :
(i) Conc. HNO
3
can be transported in aluminium container.
(ii) A mixture of dilute NaOH and aluminium pieces is used to open
drain.
(iii) Graphite is used as lubricant.
(iv) Diamond is used as an abrasive.
(v) Aluminium alloys are used to make aircraft body.
(vi) Aluminium utensils should not be kept in water overnight.
(vii) Aluminium wire is used to make transmission cables.
11.23 Explain why is there a phenomenal decrease in ionization enthalpy from
carbon to silicon ?
11.24 How would you explain the lower atomic radius of Ga as compared to Al ?
11.25 What are allotropes? Sketch the structure of two allotropes of carbon namely
diamond and graphite. What is the impact of structure on physical
properties of two allotropes?
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333THE p-BLOCK ELEMENTS
11.26 (a) Classify following oxides as neutral, acidic, basic or amphoteric:
CO, B
2
O
3
, SiO
2
, CO
2
, Al
2
O
3
, PbO
2
, Tl
2
O
3
(b) Write suitable chemical equations to show their nature.
11.27 In some of the reactions thallium resembles aluminium, whereas in others
it resembles with group I metals. Support this statement by giving some
evidences.
11.28 When metal X is treated with sodium hydroxide, a white precipitate (A) is
obtained, which is soluble in excess of NaOH to give soluble complex (B).
Compound (A) is soluble in dilute HCl to form compound (C). The compound
(A) when heated strongly gives (D), which is used to extract metal. Identify
(X), (A), (B), (C) and (D). Write suitable equations to support their identities.
11.29 What do you understand by (a) inert pair effect (b) allotropy and
(c) catenation?
11.30 A certain salt X, gives the following results.
(i) Its aqueous solution is alkaline to litmus.
(ii) It swells up to a glassy material Y on strong heating.
(iii) When conc. H
2
SO
4
is added to a hot solution of X,white crystal of an
acid Z separates out.
Write equations for all the above reactions and identify X, Y and Z.
11.31 Write balanced equations for:
(i) BF
3
+ LiH →
(ii) B
2
H
6
+ H
2
O →
(iii) NaH + B
2
H
6
→
(iv) H
3
BO
3
∆
→
(v) Al + NaOH →
(vi) B
2
H
6
+ NH
3
→
11.32. Give one method for industrial preparation and one for laboratory
preparation of CO and CO
2
each.
11.33 An aqueous solution of borax is
(a) neutral (b) amphoteric
(c) basic (d) acidic
11.34 Boric acid is polymeric due to
(a) its acidic nature (b) the presence of hydrogen bonds
(c) its monobasic nature (d) its geometry
11.35 The type of hybridisation of boron in diborane is
(a) sp (b) sp
2
(c) sp
3
(d) dsp
2
11.36 Thermodynamically the most stable form of carbon is
(a) diamond (b) graphite
(c) fullerenes (d) coal
11.37 Elements of group 14
(a) exhibit oxidation state of +4 only
(b) exhibit oxidation state of +2 and +4
(c) form M
2–
and M
4+
ions
(d) form M
2+
and M
4+
ions
11.38 If the starting material for the manufacture of silicones is RSiCl
3
, write the
structure of the product formed.
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