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p – BLOCK ELEMENTS
Introduction
There are six groups of p-block elements in the periodic table numbering from 13to 18. Boron,
carbon, nitrogen, oxygen, fluorine and helium head the groups. Their valence shell electronic
configuration is ns² np1-6 (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.
In groups 13, 14 and 15, the group oxidation state is the most stable state for lighter elements of the
group. However, the oxidation state two units less than the group oxidation state becomes
progressively more stable down a group. This is due to the reluctance of ns² electrons to participate
in bond formation in the case of heavier elements. This phenomenon is known as inert pair effect.
Since p-block contains non-metals (and metalloids), these elements have higher electronegativities
and higher ionisation enthalpies. In contrast to metals which form cations, non-metals readily form
anions.
The combined effect of size and availability of cf orbitals considerably influences the ability of these
elements to form p bonds. The first member of a group differs from the heavier members in its ability
to form pp -pp multiple bonds to itself ( e.g., C=C, C° C, N° N) and to other second row elements
e.g., C=0, C=N, C^ N, N=0). This type of p – bonding is not particularly strong for the heavier p-
block elements. The heavier elements do form p bonds but this involves d orbitals.
Group 13 Elements: The Boron Family
Electronic Configuration
The outer electronic configuration of these elements is ns² np¹. This difference in electronic structures
affects the other properties and consequently the chemistry of all the elements of this group.
Atomic Radii
On moving down the group, 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 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).
Ionization Enthalpy
The ionisation enthalpy values as expected from the general trends do not decrease 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.
Electronegativity
Down the group, electronegativity first decreases from B to Al and then increases marginally.
Physical Properties
Boron is non-metallic in nature. It is extremely hard and black coloured solid. It exists in many
allotropic forms.
Some Important Compounds Of Boron
Borax
It is the most important compound of boron. Formula of the compound is Na 2
B
4
O
7
.10H
2
O. In fact
it contains the tetranuclear units
2
(^4 5 )
B O OH
and correct formula; therefore, is
Na 2
[B
4
O
5
(OH)
4
].8H
2
O.
On heating, borax first loses water molecules. On further heating it turns into a transparent liquid,
which solidifies into glass like material known as borax bead.
Orthoboric acid
Orthoboric acid, H 3
B
3
is a white crystalline solid, with soapy touch. It is sparingly soluble in water
but highly soluble in hot water.
Na 2
B
4
O
7
O 2NaCl + 4B(OH) 3
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+
Structure of boric acid is given below.
Diborane (B 2
H
6
The simplest boron hydride is diborane (B 2
H
6
). Diborane can be prepared by treating BF3 with
lithium aluminium hydride in ether. A convenient laboratory method is oxidation of sodium borohydride
with iodine. 3 4 2 6 3 4BF 4LiAlH 2B H 3LiF 3AlF
2NaBH 4
B
2
H
6
On a commercial scale, diborane is produced by the action of BF3 on sodium hydride.
3 2 6 450 K
BF 6NaH B H 6NaF
Diborane is a colourless toxic gas. It catches fire on exposure to air releasing large amount of
energy.
B
2
H
6
+ 6H
2
O 2B(OH)
3
+ 6H
2
Reaction of diborane with NH 3
gives an addition product B 2
H
6
.2NH
3
which on heating gives borazine
(B
3
N
3
H
3
), commonly known as inorganic benzene due to its structural similarity with benzene.
Boron forms a series of hydridoborates, the most important being (BH 4
(sodium
borohydride) is a good reducing agent.
Each boron atom in B 2
H
6
is sp³ hybridised. The structure contains two types of H- atoms the four-
terminal hydrogen atoms and two bridged hydrogen atoms. The four-terminal H atoms and two B
atoms lie in the same plane. Above and below this plane lie the bridged H atoms. B-H bonds
formed by the terminal hydrogen atoms are normal covalent bonds while the bridge B-H bonds are
three centre two-electron bonds. Each B atom forms four bonds even though boron has only three
valence electrons. Hence B 2
H
6
is an electron deficient compound.
Group 14 Elements:
The Carbon Family
Carbon, silicon, germanium, tin, and lead form the carbon family.
Occurrence:
Carbon is widely distributed in nature in the free and combined states. Graphite, diamond, coal, etc
are elemental forms of carbon while in the combined state it occurs as metal carbonates, hydrocarbons
and CO 2
in air. Silicon is present in nature as silica and silicates. Ge is found only in traces. Tin
occurs as cassiterite (SnO 2
) and lead as galena (PbS)
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 (i) Reactivity towards oxygen
All members when heated in oxygen form oxides. There are mainly two types of oxides, monoxide,
and dioxide of formula MO and MOs respectively.
(ii) Reactivity towards water
2 2 2 Sn 2H O SnO 2H
(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.
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:
I mportant Trends And Anomalous Behaviour Of Carbon
Carbon 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.
Carbon has the ability to form pp – pp multiple bonds with itself and with other atoms of small size
and high electronegativity.
Few examples are: C=C, C C, C=0, C=S, and C ^ N. Carbon atoms have the tendency to link
with one another through covalent bonds to form chains and rings. This property is called catenation.
Allotropes Of Carbon
Diamond
3
o f diamond
hybridisation sp
H C 1.90KJ/mol
It has a crystalline lattice. In diamond, each carbon atom undergoes sp³ hybridisation and linked to
four other carbon atoms by using hybridised orbitals in tetrahedral fashion. The C-C bond length is
154 pm. In this structure, 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. (1 carat diamond = 200 mg)
Graphite
2
o f graphite
hybridisation sp
H C 0
Graphite has layered structure. 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 sp²
hybridisation and makes three sigma bonds with three neighbouring carbon atoms. Fourth electron
forms a p 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.
Fullerenes
Fullerenes are prepared by heating graphite in an electric arc in the presence of helium or argon.
The sooty material formed by condensation of the vapours consists of C 60
with smaller amounts of
C
70
and other fullerenes. C 60
is named as Buckminster fullerence. The general name fullerence
refers to the family of spheroidal carbon-cage molecules. The shape of C 60
resembles that of a
soccer ball. It contains twelve five-membered rings and twenty 6-membered rings of carbon. The
6-membered rings are fused both to other five and six membered rings. However, the 5-membered
rings are fused only to six-membered rings. Both carbon-carbon single (1.435 Å) and double
(1.383 Å) bonds are present in this structure. Carbon black, coke and charcoal are impure amorphous
forms of graphite or fullerenes. Carbon black is formed by burning hydrocarbon in limited supply of
air. Charcoal and coke are obtained by heating wood and coal respectively in the absence of air.
from carrying oxygen round the body and ultimately resulting in death.
Carbon Dioxide
It is prepared by complete combustion of carbon and carbon-containing fuels in excess of air.
On commercial scale it is obtained by heating limestone. 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:
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.
Resonance structures of carbon dioxide
Silicon Dioxide, SiO 2
Quartz, cristobatite and tridymite are some of the crystalline forms of silica, and they are
interconvertible at suitable temperature. In Silicon dioxide, 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
2 2 3 2
2 4 2
SiO 2NaOH Na SiO H O
SiO 4HF SiF 2H O
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, RnSiCl (4-n)
where R is alkyl or aryl group.
