P –
Block Elements Class 11 Notes
Elements in which distinguishing or valence electron enters
in the p – orbital are known as P – block elements. The only exceptional element
in this block is helium (atomic number=2) that has the electronic configuration
1s2. The general valence shell electronic configuration of P–block
elements is ns2 np1-6. They are placed to the
right hand side of long form of the periodic table and includes elements
belonging to group 13 to 18. They consist of metals, metalloids, non – metals
and noble gases. There are six families in the p – block elements.
Group
|
Name of family
|
Outermost electronic configuration
|
Elements
|
13
|
Boron family
|
ns2 np1
|
B, Al, Ga, In, Ti
|
14
|
Carbon family
|
ns2 np2
|
C, Si, Ge, Sn, Pb
|
15
|
Nitrogen family
|
ns2 np3
|
N, P, As, Sb, Bi
|
16
|
Chalcogens i.e. oxygen family
|
ns2 np4
|
O, S, Se, Te, Po
|
17
|
Halogens i.e. fluorine family
|
ns2 np5
|
F, Cl, Br, I, At
|
18
|
Noble gases
|
ns2 np6
|
He, Ne, Ar, Kr, Xe, Rn
|
GROUP-13 ELEMENTS: BORON FAMILY
The group – 13 elements contains five elements namely boron (B), Aluminum
(Al), gallium (Ga), indium (In) and thallium (Tl). The outer electronic
configuration of group -13 element is ns2np1.
Boron is a semi – metal while aluminum and rest of the elements are metallic in
nature.
OCCURRENCE
Boron does not occur in the Free State in the nature. It is fairly rare
element. It mainly occurs as borates, orthoboric acid, borax and kernite. It
shows two isotopic forms as 10B (19%) and 11B (81%).
Aluminium also does not occur in free state. It is forth most abundant element
in the earth crust. Important minerals of aluminium are bauxite and cryolite.
Gallium, indium and thallium are less abundant elements in nature. Thallium is
highly toxic amongst the group.
PHYSICAL PROPERTIES
1. Atomic radii
On moving down the group, atomic
radius increases due increase in extra shell of electrons. But atomic radius of
gallium is less than that of aluminium. This is because of variation in the
inner core of the electronic configuration.
2. Ionization enthalpy
On moving down the group, regular decreasing trend (expected) in the
ionization enthalpy value is not observed. It decreases from B to Al, increases
in Ga and again decreases in In then increases in Tl.
IE1 : B > Al ~ Ga > In < Tl
3. Electropositive or Metallic character
The elements of group – 13 are less electropositive than S – block
elements. The electropositive character increases from B to Al and then
decreases from Al to Tl. Aluminium is metal and is most electropositive.
4. Lewis acids
Trivalent compounds of group – 13 such as BCl3, Bf3,
AlCl3 etc., have six electrons around the central atom and thus act
as Lewis acids by accepting a lone pair of electron to achieve stable
electronic configuration. Lewis acid character decreases down the group.
5. Oxidation states
Boron shows only +3 oxidation state while other elements show +1 as well
as +3 oxidation states.
CHEMICAL REACTIVITY
1. Action of air (oxygen)
Boron is unreactive towards air in crystalline form. While amorphous
boron and aluminium metal on heating in air form B2O3 and
Al2O3 also with nitrogen at high temperature they form
nitrides. All the members of this group form oxides of the type E2O3.
4E + 3O
ฮ → 2E2O3
2E + N2 →
2EN (E = B, Al, Ga, In, Tl)
Oxides of B is acidic, Al and Ga oxides are amphoteric and those of In
and TI are basic in their properties.
2. Reactivity towards water
Pure boron does not react with water while it reacts with steam at red
heat. Aluminium decomposes boiling water evolving hydrogen. Gallium and indium
are not attacked by water.
2Al + 3H2O → Al2O3 +
3H2 ↑
3. Reactivity towards halogens
All the elements of group – 13 react with halogen to form trihalides (except
TlI3)
2E + 3X2 → 2EX3
Where X = F, Cl, Br, I and E
=B, Al, Ga, In, Tl
4. Reactivity towards acids
Boron does not reacts with non – oxidizing acids (HCl and dil. H2SO4)
but with oxidizing acids (conc. H2SO4 and HNO3)
boron forms boric acid.
B + 3HNO3 → H3BO3 +
3NO2
2B + 3H2SO4 → 2H3BO3 +
3SO2
While other elements react with any acid to form trivalent salt.
2E + 6HCl
→ 2ECl3 + 3H2
↑
2E + 3H2SO4 → E2(SO4)3 + 3H2
↑
5. Reactivity towards alkalies
Boron and aluminium react with alkalies to give borates and meta –
aluminate with the evolution of H2 gas.
2B + 6NaOH
Fuse → 2Na3BO3 + 3H2
↑
2Al + 2NaOH
+ 2H2O Fuse → 2NaAlO2 + 3H2
↑
Indium and thallium are not affected by alkalies.
6. Reducing action of boron
Boron is a powerful reducing agent and it reduces CO2 and SiO2
to C and Si respectively.
4B + 3CO2 →
2B2O3
+ 3C
4B + 3SiO2 →
2B2O3
+ 3Si
ANOMALOUS PROPERTIES OF FIRST ELEMENT BORON
Boron differs from other members of group – 13 elements because of its
small size (88pm), higher electronegativity (2.0) and higher value of
ionization enthalpy (801 KJmol-1). It shows following different
properties.
·
Boron is a non – metal while other members are metals.
·
Boron shows allotropy while other members do not.
·
Boron has highest melting and boiling point amongst the
elements of group – 13.
·
Oxides and hydroxides of boron are weakly acidic, of
aluminium are amphoteric while those of rest of the members are weakly basic.
·
Boron hydride is quite stable while hydrides of other
elements are less stable.
IMPORTANT COMPOUNDS OF BORON
1) Borax
It is also known as sodium
tetraborate decahydrate (Na2, B4O7.
10H2O) Borax occurs
naturally as tincal. It is obtained from the mineral colemanite by boiling it
with a solution of Na2CO3.
Ca2B6O11
+ 2Na2CO3 → Na2B4O7
+
2NaBO2 + 2CaCO3
Colemanite Borax
Properties
i.
Aqueous solution of boron is alkaline due to hydrolysis.
Na2B4O7
+ 7H2O
→
2NaOH
+ 4H3BO3
Strong alkali weak acid
ii.
Boron when heated with ethyl alcohol and concentrated H2SO4,
it gives volatile vapors of triethyl borate which burn with green edged flame.
This is used as test for
borate radical in quantitative analysis.
iii.
The borax bead is used to detect colored radicals. It is
transparent glassy mass.
2] Orthoboric acid (H3BO3)
Orthoboric acid is a white
crystalline solid, with a soapy touch. It is obtained from borax by treading
with dil. HCl or dil. H2SO4.
Na2B4O7
+
2HCl + 5H2O →
NaCl + 4H3BO3
It can also be obtained from mineral
colemanite by passing SO2 through mixture of powdered mineral and
boiling water.
Ca2B6O11 +
4SO2 + 11H2O →
2Ca(HSO3)2
+ 6H3BO3
Properties
i.
It is a weak monobasic acid and behaves as a Lewis acid it
accepts pair of electron from OH- ion of H2O.
H3BO3 + H2O →
[B(OH)4]-
+ H+
ii.
On heating it at 370K, it forms metaboric acid (HBO2)
which on further heating gives boric oxide (B2O3).
H3BO3 ฮ → HBO2 ฮ → B2O3
iii.
With NaOH it forms sodium metaborate.
H3BO3 +
NaOH → NaBO2 + 2H2O
iv.
With ethyl alcohol and conc. H2SO4 it
gives triethylborate which burns with green edged flame.
H3BO3 + 3C2H5OH conc. H2SO4 →
B(OC2H5)3 + 3H2O
3] Diborane (B2H6)
It is the simplest boron hydride. It
is prepared by treating boron trifluoride with LiAlH4 in diethyl
ether.
4BF3 + 3LiAH4 → 2B2H6 +
3LiF + 3AlF3
In laboratory the diboron is prepared
by the oxidation of sodium borohydride wit h
iodine.
2NaBH4 + I2 →
B2H6
+ 3NaI + H2
In industry it is prepared by the
reaction of BF3 with NaH.
2BF3 + 6NaH
→ B2H6 +
6NaF
Properties
i.
Diborone is colorless, highly toxic gas. It has boiling point
of 180K. It catches fire spontaneously when it is exposed to atmospheric air.
It burns in oxygen. The reaction is highly exothermic.
B2H6
+ 3O2 → B2O3 + 3H2O
ฮH = -1976KJmol-1
ii.
It is readily hydrolysed by water to form boric acid.
B2H6
+
6H2O → 2H3BO3 + 6H2
iii.
With lewis bases, diborane first undergoes cleavage to form borane
(BH3) which then reacts to form adducts.
B2H6
+ 2NMe3 → 2BH3.NMe3
B2H6
+ 2CO → 2BH3.Cn
iv.
It reacts with strong alkalies to form metaborates.
B2H6 + 2KOH
+ 2H2O → 2KBO2 + 6H2
v.
When diborane reacts with LiH in presence of ether it farms
lithium borohydride.
2LiH + B2H6 Ether → 2LiBH4
GROUP – 14 ELEMENTS : CARBON
FAMILY
The outer electronic configuration of
group – 14 elements is ns2np2. This group includes
carbon (C), silicon (Si), germanium (Ge), tin (Sn) and lead (Pb). C and Si are
non metals, Ge is a metalloid where Pb and Sn are metals.
OCCURRENCE
Carbon is widely distributed in
nature in free state as well as in the combined form. In elemental form it
occurs as coal, graphite, diamond and in combined form it occurs as metal
carbonates, hydrocarbons and carbon dioxide. Silicon is the second (27.7% by
mass) most abundant element in the earth crust. It is present in nature as
silicon (SiO2) and Silicates. Germanium exists only in traces. Tin
occurs as SnO2 and lead as galena (PbS).
PHYSICAL PROPERTIES
1. Atomic radii
Atomic radii of group – 14 elements
regularly increase down the group. There is large increase in atomic radius
from C to Si but from Si inwards there is small increase due to ineffective
shielding of the valence electrons by intervening d – and f – electrons.
2. Ionization enthalpy
In general ionization enthalpy
decrease down the group. From C to Si there is large decrease in ionization
enthalpy, but from Sn to Pb there is slight
increases.
3. Electronegativity
Group – 14 are slightly more
electronegative than corresponding group – 13 elements. Carbon (2.5) is most
electronegative elment in this group. Electronegativity values from Si to Pb
are almost same.
4. Oxidation states
Carbon shows only +4 oxidation state
while other elements of the group show +4 as well as +2 oxidation states.
CHEMICAL REACTIVITY
1. Reactivity towards oxygen
All elements of group – 14 when
heated in oxygen forms monoxide and dioxide.
2M
+ O2 →
2MO
M
+ O2 →
MO2
Where, M = C, Si, Ge, Sn, Pb.
The dioxides of C, Si and Ge are
acidic in nature and that of Sn and Pb are amphoteric in nature. The monoxides
of C is neutral, Ge is distinctly acidic while Sn and Pb are amphoteric.
2. Reactivity towards water
Carbon, Silicon , germanium are not
affected by water but react on red heating to form oxides. Lead is unaffected
by water. Tin decomposes steam to form dioxide and hydrogen gas.
Sn + 2H2O ฮ → SnO2 +
2H2 ↑
3. Reactivity towards halogens
Carbon forms only tetrahalide while
other elements of group react directly with halogens to form dihalides (MX2)
as well as tetrahalides (MX3).
M
+ X2 → MX2 (except carbon)
M
+ 2X2 →
MX4
Stability of dihalides increases down
the group while that of tetrahalides decreases down the group.
4. Reactivity towards alkalies
Carbon is unaffected by alkalies
while Si, Sn and Pb are Slowly attacked by cold alkali to form silicate,
stannates and plumbates respectively.
M + 2NaOH
+ H2O → Na2MO3 + 2H2
(where, M = Si, Sn, Pb)
ANOMALOUS PROPERTIES OF CARBON
Carbon differs from rest of the
element of group – 14 due to its smaller size, higher electronegativity and
high ionization enthalpy.
Some points of difference are –
·
It is the hardest element of the group – 14. It has the
highest ionization enthalpy and is most electronegative element of this group.
·
It’s maximum covalency is 4 as d – electrons are absent in
the valence shell.
·
It has maximum property of catenation. It can form multiple
bonds. As result carbon forms large numbers of compounds.
·
Carbon dioxides (CO2) is a gas while the oxides of
other elements are solids.
CATENATION
The property of self – linking of
atoms of an element to form straight or unbranched chains and rings of
different sizes by covalent bonds is called catenation. The tendency of an
element to form a chains depends upon the strength of the element – element
bond. Among the group 14 elements, C – C
bond strength (348 KJmol-1) is maximum, therefore carbon shows
maximum tendency for catenation. The order of catenation is C >> Si >
Ge ~ Sn. Lead does not show catenation.
ALLOTROPES OF CARBON
Allotropy is the phenomenon in which
on element exist in a different physical forms having different physical
properties but chemical properties are same or similar. Among the elements of
group – 14, carbon, silicon and tin have allotropic modifications. Carbon has
the largest number of allotropic modifications.
Diamond
Diamond has three dimensional
structur each carbon atom is bonded to four other carbon atoms which occupy
four corners of a tetrahedron and each of these to four more carbon atoms. In
diamond carbon atom is SP3 hydridized. C – C bond length is 154pm.
Diamond has high melting point (3873K) and high density (3.51gmCm-1).
It is non conductor of electricity and is non conductor of electricity and is
the hardest on the earth.
Graphite
Graphite has two dimensional sheet
like structure and these are fused system of planner hexagonal ring. In the
planner hexagonal rings, each carbon atom is bonded to three other carbon
atoms. In graphite carbon atom undergoes SP2 hybridization. In
graphite the distance between two successive layers is 340 pm while bond length
is 141.5pm. It conducts electricity.
Fullerenes or Bucky balls
Fullerenes is newly discovered allotrope
of carbon in 1985. Fullerenes are made by heating of graphite in an electric
arc in presence of inert gases such as helium or organ. It is cage like
molecule with chemical formula C60. It has a shape like soccer ball
and called Buckminster – fullerene. Each carbon atom is SP2
hybridized the average C – C distance is 144pm.
IMPORTANT COMPOUNDS OF CARBON :
OXIDES OF CARBON
1] Carbon monoxide (CO)
i.
It is prepared by direct combination of carbon in limited
supply of oxygen.
2C + O2 ฮ → 2CO
ii.
By laboratory method pure CO is prepared by dehydration of
formic acid with concentrated H2SO4 at 373K.
HCOOH H2SO4 373K → CO + H2O
iii.
On commercial scale it is prepared by the passage of steam
over hot coke.
C + H2O →
CO + H2
Water gas
iv.
When air is used instead of steam a mixture of CO and N2
is produced, which is called producer gas.
2C + O2
+
4N2 1273K → 2CO
+ 4N2
Properties
It is colorless and odorless gas and almost insoluble in water. It is
powerful reducing in water. It is powerful reducing agent and reduces almost
all the metal oxides. CO is used in the extraction of many metals from their
oxides.
Fe2O3
+ 3CO ฮ → 2Fe
+ 3CO2
CuO + CO
ฮ → Cu
+ CO2
ZnO + CO
→ Zn +
CO2
It is a poisonous gas, burns with blue flame and when inhaled produces
suffocation and finally death.
2] Carbon dioxide
i.
It is prepared by complete combustion of carbon and carbon
containing fuels in excess of air.
C + O2 ฮ → CO2
CH4 + 2O2 ฮ → CO2 + 2H2O
ii.
It is prepared in laboratory by the action of dilute HCl on
CaCO3.
CaCO3 +
2HCl → CaCl2 + CO2 + H2O
iii.
On commercial scale it is obtained by heating limestone.
CaCO3 1600K → CaO + CO2
iv.
It is obtained as a byproduct in manufacture of ethyl alcohol
by fermentation of glucose or fructose.
C6H12O6 zymase
→ 2C2H5OH +
2CO2
Properties
Carbon dioxide is colorless, Odorless
and tasteless gas. It is slightly soluble in water and is heavier than air. CO2
is acidic in nature it combines with alkalies to form metal carbonates.
CO2 acts as an oxidizing
agent. When heated with Zn, Fe or C these are oxidized.
Zn
+ CO2 ฮ →
ZnO
+ CO
CO2 +
C ฮ → 2CO
In photosynthesis, CO2 is
converted to glucose and oxygen.
6CO2 +
12H2O → C6H12O6 + 6H2O + 6O2
CO2 is reacts with NH3
at 453K – 473K under a pressure of 220 atmosphere to give urea as the final product.
2NH3 + CO2 453K – 473K , 220atm → NH2CONH2 + H2O
Urea
COMPOUNDS OF SILICON
Silicon dioxide (SiO2)
Silicon dioxide is commonly known as
silica. 95% of the earth crust is made up of silica and silicates. It occurs as
quartz, cristabalite, and tridymite which are crystalline.SiO2 is
covalent in nature and in normal form it is non reactive because of very high
Si – O bond enthalpy. However it is attacked by HF and NaOH.
SiO2 + 4HF
→ SiF4 + 2H2O
SiO2 +
2NaOH → Na2SiO3 + H2O
Silicon tetrachloride (SiCl4)
It is prepared by action of chlorine
on silicon
Si
+ 2Cl2 →
SiCl4
This tetrahalide is covalent in
nature and possesses tetrahedral geometry. CCl4 is not hydrolysed by
water but SiCl4 gets easily hydrolysed.
SiCl4 + 4H2O → Si(OH)4 +
4HCl
Silicones
Silicones are synthetic organosilicon
compounds containing repeated R2SiO units held by Si – O – Si
linkage. These are formed by the hydrolysis of alkyl or aryl substituted
chlorosilanes and their subsequent polymerization.
Silicates
Silicates are metal derivatives of
silicic acid, H4SiO4 or Si(OH)4.
Silicates are formed by heating metal
oxide or metal carbonates with sand.
e.g. Na2CO3 fused with sand, SiO2 → Na4SiO4
Silicates can also be termed as
metallic salts of silicon – oxygen anions. Silicates have basic unit of SiO44-,
each silicon atom is bounded with four oxide ions tetrahedrally. Rocks, clays,
soils are made up of silicates of aluminium, iron, magnesium and other metals.
