CBSE Class 11 Syllabus 2027 Chemistry Reading Material

FOR CBSE CLASS 11 SYLLABUS EXAM PREPARATION

CBSE Class 11 Syllabus
2027
Question Paper ·
Chemistry
EXAM YEAR TYPE SUBJECT

CBSE Class 11 Syllabus 2027 Question Paper Chemistry
DETAILS

Reading Material

Notes · Sample Papers · Previous Year Papers · Mock Tests

m . co
c. o m

s e
s em l a
g l a ag
READING MATERIAL
CLASS-XI (2026-27)
CHEMISTRY
Adapted from NCERT textbook of class XI 2006 (043)
edition reprinted edition pre
rationalization
m
m .co
Adapted
Adaptedfrom
fromNCERT
m .co
NCERTtextbook
textbookof
ofClass
class XI 2022-23 edition.
2006 edition.
s e m
s e
STATESlaOF MATTER g la
a
ag
INTRODUCTION
In previous units we have learnt about the properties related to single particle of matter,
such as atomic size, ionization enthalpy, electronic charge density, molecular shape and
polarity, etc. Most of the obs erva ble characteristics of chemical systems with which we

m
are familiar represent bulk properties of matter, i.e., the properties associated with a

.co
collection of a large number of atoms, ions or molecules. For example, an individual

m
molecule of a liquid does not boil but the bulk boils. Collection of water molecules have

s e
wetting properties; individual molecules do not wet. Water can exist as ice, which is a solid;

l a
it can exist as liquid; or it can exist in the gaseous state as water vapour or steam. Physical
g
a
properties of ice, water and steam are very different. In all the three states of water chemical
composition of water remains the same i.e., H 2 O. Characteristics of the three states of water
depend on the energies of molecules and on the manner in which water molecules aggregate.
Same is true for other substances also.
Chemical properties of a substance do not change with the change of its physical state; but
rate of chemical reactions do depend upon the physical state. Many times in calculations
while dealing with data of experiments we require knowledge of the state of matter.
Therefore, it becomes necessary for a chemist to know the physical laws which govern the
behaviour of matter in different states.

m
.co
THE GASEOU STATE
m
.co m
This is the si mplest state of matt er. Throughout our life we remain immersed in the ocean

s e
of air which is a mixture of gases. We spend our life in the lowermost layer of the atmosphere

s em l a
called troposphere, which is held to the surface of the earth by gravitational force. The thin layer
g
la a
of atmosphere is vital to our life. It shields us from harmful radiations and contains substances

ag like dioxygen, dinitrogen, carbon dioxide, water vapour, etc.
Let us now focus our attention on the behaviour of substances which exist in the gaseous state
under normal conditions of temperature and pressure. A look at the periodic table shows that
only eleven elements exist as gases under normal conditions (Fig5.4).

om . co
. c e m
e m las
as ag

gl For more Question Papers, Sample Papers, Notes & Syllabus visit Page 1 of 74

centuries on the physical properties of gases.
The first reliable measurement on properties
of gases was made by Anglo-Irish scientist
Robert Boyle in 1662. The law which he
formulated is known as Boyle’s Law. Later
on attempts to fly in air with the help of hot
air balloons motivated Jaccques Charles and
Joseph Lewis Gay Lussac to dis cov er
additional gas laws. Contribution from
Avogadro and others provided lot of
information about gaseous state.

Boyle’s Law (Pressure - Volume
Relationship)
Fig. 5.4 Eleven elements that exist as gases On the basis of his experiments, Robert Boyle
The gaseous state is characterized by the reached to the conclusion that at constant
following physical properties. temperature, the pressure of a fixed
 Gases are highly compressible. amount (i.e., number of moles n) of gas
varies inversely with its volume. This is
 Gases exert pressure equally in all
known as Boyle’s law. Mathematically, it can
directions.
be written as
 Gases have much lower density than the
1
p  ( at constant T and n) (5.1)
V
solids and liquids. 1
 The volume and the shape of gases are  p = k1 (5.2)
V
not fixed. These assume volume and shape
of the container. where k1 is the proportionality constant. The
 Gases mix evenly and completely in all value of constant k1 depends upon the
proportions without any mechanical aid. amount of the gas, temperature of the gas
and the units in which p and V are expressed.
Simplicity of gases is due to the fact that
the forces of interaction between their On rearranging equation (5.2) we obtain
molecules are negligible. Their behaviour is
pV = k1 (5.3)
governed by same general laws, which were
discovered as a result of their experimental It means that at constant temperature,
studies. These laws are relationships between product of pressure and volume of a fixed
measurable properties of gases. Some of these amount of gas is constant.
properties like pressure, volume, temperature If a fixed amount of gas at constant
and mass are very i mportant bec aus e temperature T occupying volume V1 at
relationships between these variables describe pressure p1 undergoes expansion, so that
state of the gas. Interdependence of these volume becomes V2 and pressure becomes p2,
variables leads to the formulation of gas laws. then according to Boyle’s law :
In the next section we will learn about gas
laws. p V = p V = constant (5.4)
1 1 2 2

THE GAS LAWS
p1 V2
The gas laws which we will study now are the  = (5.5)
result of research carried on for several p2 V1

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 2 of 74

Figure 5.5 shows two conventional ways
of grap hically presenting Boyle’s law .
Fig. 5.5 (a) is the graph of equation (5.3) at
different temperatures. The value of k1 for
each curve is different because for a given
mass of gas, it varies only with temperature.
Each curve corresponds to a different
constant temperature and is known as an
isotherm (constant temperature plot). Higher
curves correspond to higher temperature. It
should be noted that volume of the gas
doubles if pressure is halved. Table 5.1 gives
effect of pressure on volume of 0.09 mol of
CO2 at 300 K.

Fig 5.5 (b) represents the graph between
1
Fig. 5.5(a) Graph of pressure, p vs. Volume, V of p and . It is a straight line passing through
a gas at different temperatures. V
origin. However at high pressures, gases
deviate from Boyle’s law and under such
conditions a straight line is not obtained in
the graph.
Experiments of Boyle, in a quantitative
manner prove that gases are hi ghl y
compressible because when a given mass of
a gas is compressed, the same number of
molecules occupy a smaller space. This means
that gases become denser at high pressure.
A relationship can be obtained between
density and pressure of a gas by using Boyle’s
law :
By definition, density ‘d’ is related to the
mass ‘m’ and the volume ‘V’ by the relation
m
d  . If we put value of V in this equation
1 V
Fig. 5.5 (b) Graph of pressure of a gas, p vs.
V
Table 5.1 Effect of Pressure on the Volume of 0.09 mol CO2 Gas at 300 K.
Pressure/104 Pa Volume/10–3 m3 (1/V )/m–3 pV/102 Pa m3

2.0 112.0 8.90 22.40
2.5 89.2 11.2 22.30
3.5 64.2 15.6 22.47
4.0 56.3 17.7 22.50
6.0 37.4 26.7 22.44
8.0 28.1 35.6 22.48
10.0 22.4 44.6 22.40

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 3 of 74

m . co
c. o m

s e
s em l a
g l a ag

from Boyle’s law equation, we obtain the  273.15 +t 

relationship.  Vt = V0  273.15 (5.6)
 
m 
d  p = k p At this stage, we define a new scale of
k
 1  temperature such that t °C on new scale is
m
.co
This shows that at a constant
given by T = 273.15 + t and 0 °C will be given
m
temperature, pressure is directly proportional

.co m
to the density of a fixed mass of the gas. by T0 = 273.15. This new temperature scale

m
is called the Kelvin temperature scale or
s e
e la
Absolute temperature scale.

las Thus 0°C on the celsius scale is equal to
ag
ag 273.15 K at the absolute scale. Note that
degree sign is not used while writing the
temperature in absolute temperature scale,
i.e., Kelvin scale. Kelvin scale of temperature
is also called Thermodynamic scale of
temperature and is used in all scientific
works.
Thus we add 273 (more precisely 273.15)
to the celsius te mperature to obt ai n
temperature at Kelvin scale.
m
.co
If we write Tt = 273.15 + t and T0 = 273.15

e m in the equation (5.6) we obtain the

l as relationship
 Tt 
the volume of balloon should be less a
g
than
V = V
t 0  
 T 0 
V T
t
 t = (5.7)
Charles’ Law (Temperature - Volume V0 T0
Relationship) Thus we can write a general equation as
Charles and Gay Lussac performed several follows.
experiments on gases independently to V2 T
= 2 (5.8)
improve upon hot air balloon technology.

m
Their investigations showed that for a fixed V1 T1

m .co
mass of a gas at constant pressure, volume
V1 V2

.co
 
m
of a gas increases on increasing temperature

m
and decreases on cooling. They found that T1 T2
s e
s e for each degree rise in temperature, volume V
g l a
g l a of a gas increases by
1
of the original

T
= constant = k2
a (5.9)

a 273.15
volume of the gas at 0 ° C. Thus if volumes of Thus V = k2 T (5.10)
the gas at 0 °C and at t ° C are V 0 and Vt The value of constant k2 is determined by
respectively, then the pressure of the gas, its amount and the
V = V + t units in which volume V is expressed.
V
t 0 273.15 0 Equation (5.10) is the mathematical
 V t = V 0 1 t 
 expression for Charles’ law, which states that
273.15 pressure remaining constant, the volume
 

om . co
. c e m
e m las
as ag

gl For more Question Papers, Sample Papers, Notes & Syllabus visit Page 4 of 74


of a fixed mass of a gas is directly
proportional to its absolute temperature.
Charles found that for all gases, at any given
pressure, graph of volume vs temperature (in
celsius) is a straight line and on extending to
zero volume, each line intercepts the
temperature axis at – 273.15 ° C. Slopes of
lines obtained at different pressure are
different but at zero volume all the lines meet
the temperature axis at – 273.15 ° C (Fig. 5.6).





Gay Lussac’s Law (Pressure-
Temperature Relationship)
Pressure in well inflated tyres of automobiles
is almost constant, but on a hot summer day
this increases considerably and tyre may
Fig. 5.6 Volume vs Temperature ( °C) graph burst if pressure is not adjusted properly.
Each line of the volume vs temperature During winters, on a cold morning one may
graph is called isobar. find the pressure in the tyres of a vehicle
Observations of Charles can be interpreted decreased considerably. The mathematical
if we put the value of t in equation (5.6) as relationship between pressure an d
– 273.15 ° C. We can see that the volume of temperature was given by Joseph Gay Lussac
the gas at – 273.15 ° C will be zero. This means and is known as Gay Lussac’s law. It states
that gas will not exist. In fact all the gases get that at constant volume, pressure of a fixed
liquified before this temperature is reached. amount of a gas varies directly with the
The lowest hypothetical or im agi nary temperature. Mathematically,
temperature at which gases are supposed to
pT
occupy zero volume is called Absolute zero.
p
All gases obey Charles’ law at very low   constant = k
3
pressures and high temperatures. T
This relationship can be derived from
Boyle’s law and Charles’ Law. Pressure vs
temperature (Kelvin) graph at constant molar
volume is shown in Fig. 5.7. Each line of this
graph is called isochore.

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 5 of 74

will find that this is the same number which
we came across while discussing definition
of a ‘mole’ (Unit 1).
Since volume of a gas is dir ectl y
proportional to the number of moles; one mole
of each gas at standard temperature and
pressure (STP)* will have same volume.
Standard temperature and pressure means
273.15 K (0°C) temperature and 1 bar (i.e.,
exactly 10 5 pascal) pressure. Th ese
values approximate freezing temperature
of water and atmospheric pressure at sea
level. At STP molar volume of an ideal gas
or a combination of ideal gases i s
22.71098 L mol–1.
Molar volume of some gases is given in
(Table 5.2).
Table 5.2 Molar volume in litres per mole of
Fig. 5.7 Pressure vs temperature (K) graph some gases at 273.15 K and 1 bar
(Isochores) of a gas. (STP).

Argon 22.37
Avogadro Law (Volume - Amount
Relationship) Carbon dioxide 22.54
In 1811 Italian scientist Amedeo Avogadro Dinitrogen 22.69
tried to combine conclusions of Dalton’s Dioxygen 22.69
atomic theory and Gay Lussac’s law of Dihydrogen 22.72
combining volumes (Unit 1) which is now
Ideal gas 22.71
known as Avogadro law. It states that equal
volumes of all gases under the same Number of moles of a gas can be calculated
conditions of temperature and pressure as follows
contain equal number of molecules. This m
means that as long as the temperature and n= (5.12)
pressure remain constant, the volume M
depends upon number of molecules of the gas Where m = mass of the gas under
or in other words amount of the gas . investigation and M = molar mass
Mathematically we can write Thus,
V  n where n is the number of m
moles of the gas. V = k4 (5.13)
M
 V  k4 n (5.11) Equation (5.13) can be rearranged as
follows :
The number of molecules in one mole of a
gas has been determined to be 6.022 10 23 m
M = k4 = k4d (5.14)
and is known as Avogadro constant. You V

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 6 of 74

m . co
c. o m

s e
s em l a
g l a ag

Here ‘d’ is the density of the gas. We can Equation (5.18) shows that the value of R
conclude from equation (5.14) that the density depends upon units in which p, V and T are
of a gas is directly proportional to its molar measured. If three variables in this equation
mass. are known, fourth can be calculated. From
A gas that follows Boyle’s law, Charles’
m
this equation we can see that at constant

.co
law and Avogadro law strictly is called an temperature and pressure n moles of any gas
m
.co m
ideal gas. Such a gas is hypothetical. It is

m
assumed that intermolecular forces are not will have the same volume because V =
nRT
s e
s e
present between the molecules of an ideal gas. p
g la
la
and n,R,T and p are constant. This equation
Real gases follow these laws only under
g will be applicable to any gas, under those a
a
certain specific conditions when forces of
interaction are practically negligible. In all
conditions when behaviour of the gas
approaches ideal behaviour. Volume of
other situations these deviate from ideal one mole of an ideal gas under ST P
behaviour. You will learn about the deviations conditions (273.15 K and 1 bar pressure) is
later in this unit. 22.710981 L mol–1. Value of R for one mole of
an ideal gas can be calculated under these
IDEAL GAS EQUATION conditions as follows :
The three laws which we have learnt till now
can be combined together in a single equation
10 Pa22.71 ×10 m 
5 –3 3

R =
which is known as ideal gas equation.
o m 1mol273.15 K 
. c= 8.314 Pa m K mol
m
3 –1 –1

s e
g l a = 8.314 10–2 bar L K –1 mol–1

At constant T and n; V 
1 a = 8.314 J K –1 mol–1
At STP conditions used earlier
Boyle’s Law
p (0 °C and 1 atm pressure), value of R is
At constant p and n; V  T Charles’ Law 8.20578 10–2 L atm K–1 mol–1.
At constant p and T ; V  n Avogadro Law Ideal gas equation is a relation between
Thus, four variables and it describes the state of
nT any gas, therefore, it is also called equation
V (5.15) of state.
m
p

.co
nT

om
Let us now go back to the ideal gas
 V = R (5.16)
c.where p
m
equation. This is the relationship for the
e
em rearranging the equation (5.16) we obtain
R is proportionality constant. On
l as
simultaneous variation of the variables. If

las g
temperature, volume and pressure of a fixed
a
g
amount of gas vary from T1, V1 and p1 to T2,

a
pV = n RT (5.17)
V2 and p2 then we can write
R = pV
 (5.18) pV p V
1 1 2 2
nT = nR and = nR
R is called gas constant. It is same for all T1 T2
gases. Therefore it is also called Universal p1V1 pV
Gas Constant. Equation (5.17) is called ideal  = 2 2
T1 T2 (5.19
gas equation

om . co
c m
Equation (5.19) is a very useful equation. If out of six, values of five variables are known, the value of
.
unknown variable can be calculated from the equation (5.19). This equation is also known as
m s e
e la
Combined gas law.

as ag

gl For more Question Papers, Sample Papers, Notes & Syllabus visit Page 7 of 74

5.6.2 Dalton’s Law of Partial Pressures

The law was formulated by John Dalton in
1801. It states that the total pressure
exerted by the mixture of non-reactive
gases is equal to the sum of the partial
pressures of individual gases i.e., the
pressures which these gases would exert if
they were enclosed separately in the same
volume and under the same conditions of
temperature. In a mixture of gases, the
pressure exerted by the individual gas is
called partial pressure. Mathematically,
p = p +p +p +......(at constant T, V) (5.23)
Total 1 2 3


640 mL  298 K 

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 8 of 74

n2RT
p  (5.26)
2 V

n3RT
p  (5.27)
3 V
where n1 n2 and n3 are number of moles of
these gases. Thus, expression for total
pressure will be
pTotal = p1 + p2 + p3

= n RT  n RT  n RT
1 2 3
V V V
RT
= (n + n + n ) (5.28)
1 2 3
V
On dividing p1 by ptotal we get
  RTV
=  n +n 1 +n  RTV
p1 n
p
total  1 2 3 


n n
= n +n1 +n = n1 = x1
1 2 3

where n = n1+n2+n3
x1 is called mole fraction of first gas.
Thus, p1 = x1 ptotal
Similarly for other two gases we can write
p2 = x2 ptotal and p3 = x3 ptotal
Thus a general equation can be written as
pi = xi ptotal (5.29)
where pi and xi are partial pressure and mole
fraction of ith gas respectively. If total pressure KINETIC MOLECULAR THEORY OF GASES
of a mixture of gases is known, the equation So far we have learnt the laws (e.g., Boyle’s
(5.29) can be used to find out pressure exerted law, Charles’ law etc.) which are concise
by individual gases. statements of experimental facts observed in
the laboratory by the scientists. Conducting
careful experiments is an important aspect
of scientific method and it tells us how the
particular system is behaving under different
conditions. However, once the experimental
facts are established, a scientist is curious to
know why the system is behaving in that way.
For example, gas laws help us to predict that
pressure increases when we compress gases

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 9 of 74

m . co
c. o m

s e
s em l a
g l a ag

but we would like to know what happens at motion of molecules will stop and gases
molecular level when a gas is compressed ? A will settle down. This is contrary to what
theory is constructed to answer such is actually observed.
questions. A theory is a model (i.e., a mental • At any particular time, different particles

m
picture) that enables us to better understand in the gas have different speeds and hence

.co
our observations. The theory that attempts different kinetic energies. Thi s
m
.co
to elucidate the behaviour of gases is known
m
assumption is reasonable because as the
as kinetic molecular theory.
m
particles collide, we expect their speed to
s e
e la
change. Even if initial speed of all the
s
Assumptions or postulates of the kinetic-

la
molecular theory of gases are given below. particles was same, the mol ec ul ar
ag
ag
These postulates are related to atoms and
molecules which cannot be seen, hence it is
collisions will disrupt this uniformity.
Consequently the particles must have
said to provide a microscopic model of gases. different speeds, which go on changing
constantly. It is possible to show that
• Gases consist of large number of identical
though the indivi dual speeds are
particles (atoms or molecules) that are so
changing, the distribution of speeds
small and so far apart on the average that
remains constant at a part ic ul ar
the actual volume of the molecules is
temperature.
negligible in comparison to the empty
space between them. They are considered • If a molecule has variable speed, then it

m
as point masses. This ass umpt io n must have a variable kinetic energy.

.co
explains the great compressibility of gases. Under these circumstances, we can talk
only about average kinetic energy. In
• There is no force of attraction between the
e m
s
kinetic theory it is assumed that average
particles of a gas at ordinary temperature
and pressure. The support for this
g l a kinetic energy of the gas molecules is

a
directly proportional to the absolute
assumption comes from the fact that
temperature. It is seen that on heating a
gases expand and occupy all the space
gas at constant volume, the pressure
available to them.
increases. On heating the gas, kinetic
• Particles of a gas are always in constant energy of the particles increases and these
and random motion. If the particles were strike the walls of the container more
at rest and occupied fixed positions, then frequently thus exerting more pressure.
a gas would have had a fixed shape which
Kinetic theory of gases allows us to derive
is not observed.
theoretically, all the gas laws studied in the
•
m
Particles of a gas move in all possible previous sections. Calculations a n d

.co
directions in straight lines. During their
m
predictions based on kinetic theory of gases

.co
random motion, they collide with each

em
agree very well with the ex peri me ntal

s
other and with the walls of the container. observations and thus establish t he

s em Pressure is exerted by the gas as a result
l a
correctness of this model.
g
la BEHAVIOUR OFaREAL GASE S:
of collision of the particles with the walls

ag •
of the container.
Collisions of gas molecules are perfectly DEVIATION FROM IDEAL GA S
elastic. This means that total energy of BEHAVIOUR
molecules before and after the collision Our theoritical model of gases corresponds
remains same. There may be exchange of very well with the experimental observations.
energy between colliding molecules, their Difficulty arises when we try to test how far
individual energies may change, but the the relation pV = nRT reproduce actual
sum of their energies remains constant. pressure-volume-temperature relationship of
If there were loss of kinetic energy, the gases. To test this point we plot pV vs p plot

om . co
. c e m
e m las
as ag

gl For more Question Papers, Sample Papers, Notes & Syllabus visit Page 10 of 74

of gases because at constant temperature, pV theoretically calculated from Boyle’s law (ideal
will be constant (Boyle’s law) and pV vs p gas) should coincide. Fig 5.9 shows these
graph at all pressures will be a straight line plots. It is apparent that at very high pressure
parallel to x-axis. Fig. 5.8 shows such a plot the measured volume is more than the
constructed from actual data for several gases calculated volume. At low pr ess ure s,
at 273 K. measured and calculated volumes approach
each other.

Fig. 5.8 Plot of pV vs p for real gas and Fig. 5.9 Plot of pressure vs volume for real gas
ideal gas and ideal gas
It can be seen easily that at constant It is found that real gases do not follow,
temperature pV vs p plot for real gases is not Boyle’s law, Charles law and Avogadro law
a straight line. There is a significant deviation perfectly under all conditions. Now two
from ideal behaviour. Two types of curves are questions arise.
seen.In the curves for dihydrogen and helium, (i) Why do gases deviate from the ideal
as the pressure increases the value of pV also behaviour?
increases. The second type of plot is seen in
(ii) What are the conditions under which
the case of other gases like carbon monoxide
gases deviate from ideality?
and methane. In these plots first there is a
negative deviation from ideal behaviour, the We get the answer of the first question if
pV value decreases with increase in pressure we look into postulates of kinetic theory once
and reaches to a mini mum val ue again. We find that two assumptions of the
characteristic of a gas. After that pV value kinetic theory do not hold good. These are
starts increasing. The curve then crosses the (a) There is no force of attraction between the
line for ideal gas and after that shows positive molecules of a gas.
deviation continuously. It is thus, found that (b) Volume of the molecules of a gas is
real gases do not follow ideal gas equation negligibly small in comparison to the space
perfectly under all conditions. occupied by the gas.
Deviation from ideal behaviour also If assumption (a) is correct, the gas will
becomes apparent when pressure vs volume never liquify. However, we know that gases
plot is drawn. The pressure vs volume plot of do liquify when cooled and compressed. Also,
experi mental data (real gas) and that liquids formed are very difficult to compress.

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 11 of 74

This means that forces of repulsion are van der Waals constants and their value
powerful enough and prevent squashing of depends on the characteristic of a gas. Value
molecules in tiny volume. If assumption (b) of ‘a’ is measure of ma gnitude of
is correct, the pressure vs volume graph of intermolecular attractive forces within the gas
experi mental data (real gas) and t hat and is independent of temperature and
theoritically calculated from Boyles law (ideal pressure.
gas) should coincide. Also, at very low te mp erat ur e,
Real gases show deviations from ideal gas intermolecular forces become significant. As
law because molecules interact with each the molecules travel with low average speed,
other. At high pressures molecules of gases these can be captured by one another due to
are very close to each other. Molecular attractive forces. Real gases show ideal
interactions start operating. At high pressure, behaviour when conditions of temperature
molecules do not strike the walls of the and pressure are such that the intermolecular
container with full impact because these are forces are practically negligible. The real gases
dragged back by other molecules due to show ideal be havio ur when press ur e
molecular attractive forces. This affects the approaches zero.
pressure exerted by the molecules on the walls The deviation from ideal behaviour can
of the container. Thus, the pressure exerted be measured in terms of compressibility
by the gas is lower than the pressure exerted factor Z, which is the ratio of product pV and
by the ideal gas. nRT. Mathematically
an 2 pV
pideal  preal  (5.30) Z = (5.32)
V2 n RT
observed correction For ideal gas Z = 1 at all temperatures
pressure term and pressures because pV = n RT. The graph
Here, a is a constant. of Z vs p will be a straight line parallel to
Repulsive forces also become significant. pressure axis (Fig. 5.10). For gases which
Repulsive interactions are short-range deviate from ideality, value of Z deviates from
interactions and are significant w he n unity. At very low pressures all gases shown
molecules are almost in contact. This is the
situation at high pressure. The repulsive
forces cause the molecules to behave as small
but impenetrable spheres. The volume
occupied by the molecules also becomes
significant because instead of moving in
volume V, these are now restricted to volume
(V–nb) where nb is approximately the total
volume occupied by the mol ec ul es
themselves. Here, b is a constant. Having
taken into account the corrections for
pressure and volume, we can rewrite equation
(5.17) as
 an2 
 p  2  V  nb  nRT (5.31)
V
 
Equation (5.31) is known as van der Waals
equation. In this equation n is number of Fig. 5.10 Variation of compressibility factor for
moles of the gas. Constants a and b are called some gases

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 12 of 74

m . co
c. o m

s e
s em l a
g l a ag

have Z  and behave as ideal gas. At high gaseous state and liquid state and that liquids
pressure all the gases have Z > 1. These are may be considered as continuation of gas
more difficult to compress. At intermediate phase into a region of small volumes and very
pressures, most gases have Z < 1. Thus gases high molecular attraction. We will also see
show ideal behaviour when the volume how we can use isotherms of gases for
m
.co
occupied is large so that the volume of the predicting the conditions for liquifaction of
m
.co
molecules can be neglected in comparison gases.
to it. In other words, the behaviour of the
e m
e m
gas becomes more ideal when pressure is very
5.9 LIQUIFACTION OF GASES

las
las
low. Upto what pressure a gas will follow the First complete data on pressure - volume -
temperature relations of a substance in both
ag
g
ideal gas law, depends upon nature of the
a
gas and its temperature. The temperature at
which a real gas obeys ideal gas law over an
gaseous and liquid state was obtained by
Thomas Andrews on carbon dioxide. He
appreciable range of pressure is called Boyle plotted isotherms of carbon dioxide at various
temperature or Boyle point. Boyle point of temperatures (Fig. 5.11). Later on it was found
a gas depends upon its nature. Above their that real gases behave in the same manner
Boyle point, real gases show pos itiv e as carbon dioxide. Andrews noticed that at
deviations from ideality and Z values are high temperatures isotherms look like that
greater than one. The forces of attraction of an ideal gas and the gas cannot be liquified
between the molecules are very feeble. Below even at very high pressure. As the

m
Boyle temperature real gases first show temperature is lowered, shape of the curve

.co
decrease in Z value with increasing pressure, changes and data shows considerable
deviation from ideal behaviour. At 30.98 °C
which reaches a minimum value. On further
e m
increase in pressure, the value of Z increases
continuously. Above explanation shows that
l as
at low pressure and high temperature gases
show ideal behaviour. These conditions are ag
different for different gases.
More insight is obtaine d in the
significance of Z if we note the following
derivation
pVreal
Z = (5.33)
n RT

m
If the gas shows ideal behaviour then

m
Videal =
n RT

c.inoequation (5.33) we have
nRT
p . On putting this value of p
m.co
m s e
s e g l a
g l a Z =
Vreal
a
a Videal (5.34)

From equation (5.34) we can see that
compressibility factor is the ratio of actual
molar volume of a gas to the molar volume of
it, if it were an ideal gas at that temperature
and pressure.
In the following sections we will see that Fig. 5.11 Isotherms of carbon dioxide at various
it is not possible to distinguish between temperatures

om . co
. c e m
e m las
as ag

gl For more Question Papers, Sample Papers, Notes & Syllabus visit Page 13 of 74

Adapted from NCERT textbook of class XI 2006 edition reprinted edition pre
rationalization

THE s -BLOCK ELEMENTS
The s-block elements of the Periodic Table are those in which the last electron enters
the outermost s-orbital. As the s-orbital can accommodate only two electrons, two groups (1 &
2) belong to the s-block of the Periodic Table. Group 1 of the Periodic Table consists of the
elements: lithium, sodium, potassium, rubidium, caesium and francium. They are
collectively known as the alkali metals. These are so called because they form hydroxides on
reaction with water which are strongly alkaline in nature. The elements of Group 2 include
beryllium, magnesium, calcium, strontium, barium and radium. These elements with the
exception of beryllium are commonly known as the alkaline earth metals. These are so called
because their oxides and hydroxides are alkaline in nature and these metal oxides are found
in the earth’s crust.
Among the alkali metals sodium and potassium are abundant and lithium, rubidium and
caesium have much lower abundances (Table 10.1). Francium is highly radioactive; its longest-
lived isotope 223Fr has a half-life of only 21 minutes. Of the alkaline earth metals calcium and
magnesium rank fifth and sixth in abundance respectively in the earth’s crust. Strontium and
barium have much lower abundances. Beryllium is rare and radium is the rarest of all
comprising only 10 –10 per cent of igneous rocks† (Table 10.2).
The general electronic configuration of s-block elements is [noble gas] ns1 for alkali metals and
[noble gas] ns2 for alkaline earth metals.

Lithium and beryllium, the first elements of Group 1 and Group 2 respectively exhibit some
properties which are different from those of the other members of the respective group. In these
anomalous properties they resemble the second element of the following group. Thus, lithium
shows similarities to magnesium and beryllium to aluminium in many of their properties. This
type of diagonal similarity is commonly referred to as diagonal relationshipin the periodic table.
The diagonal relationship is due to the similarity in ionic sizes and /or charge/ra dius ratio of
the ele me nts. Monovalent sodium and potassium ions and divalent magnesium and calcium
ions are found in large proportions in biological fluids. These ions perform important biological
functions such as maintenance of ion balance and nerve impulse conduction.
10.1 GROUP 1 ELEME NTS: AL KALI METALS
The alkali metals show regular trends in their physical and chemical properties with the
increasing atomic number. The atomic, physical and chemical properties of alkali metals are
discussed below.
10.1.1 Electronic Configuration
All the alkali metals have one valence Element Symbol Electronic configuration
electron, ns1 (Table 10.1) outside the noble Lithium Li 1s22s1
gas core. The loosely held s-electron in the
Sodium Na 1s22s22p63s1
outermost valence shell of these elements
makes them themost electropositive metals. Potassium K 1s22s22p63s23p64s1
They readily loseelectron to give monovalent Rubidium Rb 1s22s22p63s23p63d104s24p65s1
M+ ions. Hence they are never found in free Caesium Cs 1s22s22p63s23p63d104s2 4p64d105s25p66s1
state in nature. or [Xe] 6s1
Francium Fr [Rn]7s1

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 14 of 74

10.1.2 Atomic and Ionic Radii
The alkali metal atoms have the largest sizes in a particular period of the periodic table. With
increase in atomic number, the atom becomeslarger. The monovalent ions (M+) are smaller than
the parent atom. The atomic and ionic radii of alkali metals increase on moving down the group
i.e., they increase in size while going from Li to Cs.

10.1.3 Ionization Enthalpy
The ionization enthalpies of the alkali metals are considerably low and decrease down the group
from Li to Cs. This is because the effect of increasing size outweighs the increasing nuclear
charge, and the outermost electron isvery well screened from the nuclear charge.
10.1.4 Hydration Enthalpy
The hydration enthalpies of alkali metal ions decrease with increase in ionic sizes.
Li+> Na+ > K+ > Rb+ > Cs+
Li+ has maximum degree of hydration and for this reason lithium salts are mostly hydrated,
e.g., LiCl· 2H2 O

Table 10.1 Atomic and Physical Properties of the Alkali Metals

Property Lithium Sodium Potassium Rubidium Caesium Francium
Li Na K Rb Cs Fr
Atomic number 3 11 19 37 55 87
Atomic mass (g mol–1) 6.94 22.99 39.10 85.47 132.91 (223)
Electronic configuration [He] [Ne] 3s1 [Ar] [Kr] [Xe] 6s1 [Rn]
2s1 4s1 5s1 7s1
Ionization 520 496 419 403 376 ~375
enthalpy / kJ mol–1
Hydration enthalpy/kJ mol–1 –506 –406 –330 –310 –276 –
Metallic radius /pm 152 186 227 248 265 –

Ionic radiusM+/pm 76 102 138 152 167 (180)

m.p. / K 454 371 336 312 302 –
b.p / K 1615 1156 1032 961 944 –
Density / g cm –3 0.53 0.97 0.86 1.53 1.90 –
Standard potentialsE0/ V for –3.04 –2.714 –2.925 –2.930 –2.927 –
(M+ / M)
Occurrence in lithosphere† 18* 2.27** 1.84** 78-12* 2-6* ~ 10–18 *

*ppm (part per million), ** percentage by weight; † Lithosphere: The Earth’s outer layer: its crust and
part of the upper mantle

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 15 of 74

m . co
c. o m

s e
s em l a
g l a ag

10.1.5 Physical Properties
All the alkali metals are silvery white, soft and light metals. Because of the large size, these
elements have low density which increases down the group from Li to Cs. However, potassium
m
m .co
is lighter than sodium. The melting and boiling points of the alkali metals are low indicating weak

.co m
metallic bonding due to the presence of only a single valence electron in them. The alkali metals

s e
em
and their salts impart characteristic colour to an oxidizing flame. This is because the heat from the

s
flame excites the outermost orbital electron to a higher energy level. When the excited electron
a Li Na g la
Metal gl
comes back to the ground state, there is emission of radiationin the visible region as given below:
a
a Crimson red Yellow
Colour
K
Violet Red violet
Rb Cs
Blue

/nm 670.8 589.2 766.5 780.0 455.5

Alkali metals can therefore, be detected by the respective flame tests and can be
determined by flame photometry or atomic absorption spectroscopy. These elements when
irradiated with light, the light energy absorbed may be sufficient to make an atom lose electron. This
property makes caesium and potassium useful as electrodes in photoelectric cells.
m
m .co
e
10.1.6 Chemical Properties

l as
The alkali metals are highly reactive due to their large size and low ionization enthalpy. The

ag
reactivity of these metals increases down the group.
(i) Reactivity towards air: The alkali metals tarnish in dry air due to the formation of their
oxides which in turn react with moisture to form hydroxides. They burn vigorously in
oxygen forming oxides. Lithium forms monoxide, sodium forms peroxide, the other
metals for m superoxides. The superoxide O 2 – ion is stable only in the presence of large
cations such as K, Rb, Cs.

2

4 Li + O2  2 Li2O (oxide)
m
m
2 Na  O2  Na2 O2 (peroxide)
c. oM  O  MO (superoxide) m.co
m
2 2

s e
s e (M = K, Rb, Cs)
g l a
g l a a
a

om . co
. c e m
e m las
as ag

gl For more Question Papers, Sample Papers, Notes & Syllabus visit Page 16 of 74

In all these oxides the oxidation state of the alkali metal is +1. Lithium shows exceptional
behaviour in reacting directly with nitrogen of air to form the nitride, Li 3N as well. Because of their
high reactivity towards air and water, alkali metals are normally kept in kerosene oil.
(ii) Reactivity towards water: The alkali metals react with water to form hydroxide and
dihydrogen.
2M + 2H2O  2 M  2 OH  H2 (M = an alkali metal)
It may be noted that although lithium has most negative E 0 value (Table 10.1), its reaction
with water is less vigorous than that of sodium which has the least negativeE0 value among
the alkali metals. This behaviour of lithium is attributed to its small size and very high
hydration energy. Other metals of the group react explosively with water.
They also react with proton donors such as alcohol, gaseous ammonia and alkynes.

(iii) Reactivity towards dihydrogen: The alkali metals react with dihydrogen at about 673K
(lithium at 1073K) to form hydrides. All the alkali metal hydrides are ionic solids with high
melting points.
2M  H2 

(iv) Reactivity towards halogens: The alkali metals readily react vigorously with halogens
to form ionic halides, M+X–. However, lithium halides are somewhat covalent. It is
because of the high polarisation capability of lithium ion (The distortion of electron
cloud of the anion by the cation is called polarisation). The Li+ ion is very small in size and
has high tendency to distort electron cloud around the negative halide ion. Since anion
with large size can be easily distorted, among halides,lithium iodide is the most covalent
in nature.

(v) Reducing nature: The alkali metals are strong reducing agents, lithium being the most and
sodium the least powerful (Table 10.1). The standard electrode potential (E0) which
measures the reducing power represents the overall change :
M(s)M(g) sublimation enthalpy
 
M(g)M (g) e ionization enthalpy

M (g) H2OM (aq) hydration enthalpy
With the small size of its ion, lithium has the highest hydration enthalpy which accounts for
its high negative E0 value and its high reducing power.

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 17 of 74

The order is Li > Na > I– > Ag > Cl–

(vi) Solutions in liquid ammonia: The alkali metals dissolve in liquid ammonia giving deep blue
solutions which are conducting in nature.
M(x  y) NH3 [M(NH3 )x ] [e(NH3 )y ]

The blue colour of the solution is due to the ammoniated electron which absorbs energy in the
visible region of light and thus imparts blue colour to the solution. The solutions are
paramagnetic and on standing slowly liberate hydrogen resulting in the formation of amide.
M (am)  e  NH3 (1)  MNH2(am)  ½H2 (where ‘am’ denotes solution in ammonia.)
In concentrated solution, the blue colour changes to bronze colour and becomes
diamagnetic.

10.1.7 Uses
Lithium metal is used to make useful alloys, for example with lead to make ‘white metal’ bearings
for motor engines, with aluminium to make aircraft parts, and with magnesium to make ar mour
plates. It is used in thermonuclear reactions. Lithium is also used to make electrochemical cells.
Sodium is used to make a Na/Pb alloy needed to make PbEt 4 and PbMe4. These organolead
compounds were earlier used as anti-knock additives to petrol,but nowadays vehicles use lead-free
petrol. Liquid sodium metal is used as a coolant in fast breeder nuclear reactors. Potassium has a
vital role in biological systems. Potassium chloride is used as a fertilizer. Potassium hydroxide is used
in the manufacture of soft soap. It is also used as an excellent absorbent of carbon dioxide. Caesium is used
in devising photoelectric cells

10.2 ANOMOLOUS BEHAVIOUR OF LITHIUM
The anomalous behaviour of lithium is due to the : (i) exceptionally small size of its atom and ion, and (ii)
high polarising power (i.e., charge/ radius ratio). As a result, there is increased covalent character of lithium
compounds which is responsible for their solubility in organic solvents.

10.2.1 Points of Difference between Lithium and other alkali metals
i. Lithium is much harder. Its m.p. and b.p. are higher than the other alkali metals.
ii. Lithium is least reactive but the strongest reducing agent among all the alkali metals. On
combustion in air it forms mainly monoxide, Li2O and the nitride, Li3N unlike other alkali metals.
iii. LiCl is deliquescent and crystallises as a hydrate, LiCl.2H2O whereas other alkali metal chlorides do
not form hydrates.
iv. Lithium hydrogencarbonate is not obtained in the solid form while all other elements form solid
hydrogencarbonates.
v. Lithium unlike other alkali metals forms no ethynide on reaction with ethyne.
vi. Lithium nitrate when heated gives lithium oxide, Li2O, whereas other alkali metal nitrates
decompose to give the corresponding nitrite.
vii. LiF a nd Li2O are comparatively much less soluble in water than the corresponding compounds of
other alkali metals.

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 18 of 74

m . co
c. o m

s e
s em l a
g l a ag
10.3 GROUP 2 ELEMENTS: ALKALINE Element Symbol Electronic configuration
EART H METALS
The group 2 elements comprise beryllium, Beryllium Be 1s22s2
magnesium, calcium, strontium, barium and Magnesium Mg 1s22s22p 63s2
radium. They follow alkali metals in the Calcium Ca 1s22s22p 63s23p 64s2
periodic table. These (except beryllium) are
Strontium Sr 1s22s22p63s23p63d10
known as alkaline earth metals. The first 4s24p65s2
element beryllium differ rs from the rest of the
m
.co
Barium Ba 1s22s22p63s23p63d10 4s2

m
members and shows diagonal relationship to 4p64d105s25p66s2 or [Xe]6s2

.co m
alumini um. The atomic and p hysic al
e
Radium Ra [Rn]7s2

em
properties of the alkaline earth metals are

las
s
shown in Table 10.2.

laConfiguration ag
a g
10.3.1 Electronic
These elements have two electrons in the s -orbital of the valence shell (Table 10.2). Their general
electronic configuration may be represented as [noble gas] ns2. Like alkali metals, the compounds
of these elements are also predominantly ionic
10.3.2 Atomic and Ionic Radii
The atomic and ionic radii of the alkaline earth metals are smaller than those of the
corresponding alkali metals in the same periods. This is due to the increased nuclear charge in
these elements. Within the group, the atomic and ionic radii increase with increase in atomic
number.
o m
c
. Earth Metals
m
Table 10.2 Atomic and Physical Properties of the Alkaline
eCalcium Strontium
as
Mggl
Property Beryllium Magnesium Barium Radium

Atomic number 4
Be
12
a 20
Ca
38
Sr Ba
56
Ra
88
Atomic mass (g mol ) –1 9.01 24.31 40.08 87.62 137.33 226.03
2 2 2 2 2
Electronic [He] 2s [Ne] 3s [Ar] 4s [Kr] 5s [Xe] 6s [Rn] 7s2
configuration
Ionization 899 737 590 549 503 509
enthalpy (I) / kJ mol–1
Ionization 1757 1450 1145 1064 965 979

m
enthalpy (II) /kJ mol–1

m
Hydration enthalpy – 2494 – 1921 –1577 – 1443 – 1305
c. o–
.co m
(kJ/mol)
Metallic 111 160 197 215
s
222
e –

s em radius / pm

g a
l135
la a
Ionic radius 31 72 100 118 148

ag M2+ / pm
m.p. / K 1560 924 1124 1062 1002 973
b.p / K 2745 1363 1767 1655 2078 (1973)
Density / g cm–3 1.84 1.74 1.55 2.63 3.59 (5.5)
Standard potential –1.97 –2.36 –2.84 –2.89 – 2.92 –2.92
E0 / V for (M2+/ M)
Occurrence in 2* 2.76** 4.6** 384* 390 * 10 –6*
lithosphere

*ppm (part per million); ** percentage by weight
om . co
. c e m
e m las
as ag

gl For more Question Papers, Sample Papers, Notes & Syllabus visit Page 19 of 74

10.3.3 Ionization Enthalpies
The alkaline earth metals have low ionization enthalpies due to fairly large size of the atoms. Since the
atomic size increases down the group, their ionization enthalpy decreases (Table 10.2). The first
ionisation enthalpies of the alkaline earth metals are higher than those of the corresponding Group 1
metals. This is due to their small size as compared to the corresponding alkali metals. It is
interesting to note that the second ionisation enthalpies of the alkaline earth metals are smaller than
those of the corresponding alkali metals.
10.3.4 Hydration Enthalpies
Like alkali metal ions, the hydration enthalpies of alkaline earth metal ions decrease
with increase in ionic size down the group.
Be2+> Mg2+ > Ca2+ > Sr2+ > Ba2+

The hydration enthalpies of alkaline earth metal ions are larger than those of alkali metal ions. Thus,
compounds of alkaline earth metalsare more extensively hydrated than those of alkali metals, e.g.,
MgCl 2 and CaCl2 exist as MgCl 2.6H 2O and CaCl 2· 6H2O while NaCl and KCl do not form such
hydrates.
10.3.5 Physical Properties
The alkaline earth metals, in general, are silverywhite, lustrous and relatively soft but harder than the
alkali metals. Beryllium and magnesium appear to be somewhat greyish. The melting and
boiling points of these metals are higher than the corresponding alkali metals due to smaller sizes. The
trend is, however, not systematic. Because of the low ionisation enthalpies, they are strongly
electropositive in nature. The electropositive character increases down the group from Be to Ba.
Calcium, strontium and barium impart characteristic brick red, crimson and apple green colours
respectively to the flame. In flame the electrons are excited to higher energy levels and when they
drop back to the ground state, energy is emitted in the form of visible light. The electrons in
beryllium and magnesium are too strongly bound to get excited by flame. Hence, these elements do
not impart any colour to the flame. The flame test for Ca, Sr and Ba is helpful in their detection in
qualitative analysisand estimation by flame photometry. The alkaline earth metals like those of alkali
metals have high electrical and thermal conductivities which are typical characteristics of metals.
10.3.6 Chemical Properties
The alkaline earth metals are less reactive thanthe alkali metals. The reactivity of these elements
increases on going down the group.
(i) Reactivity towards air and water:
Beryllium and magnesium are kinetically inert to oxygen and water because of the formation of an
oxide film on their surface. However, powdered b e r y l l i u m b u r n s brilliantly on ignition in
air to give BeO and Be 3 N 2 . Magnesium is more electropositive and burns with dazzling
brilliance in air to give MgO and Mg3N2 . Calcium, strontium and barium are readily attacked by
air to form the oxide and nitride. They also react with water with increasing vigour even in
cold to form hydroxides.
(ii) Reactivity towards the halogens: All the alkaline earth metals combine with halogenat
elevated temperatures forming their halides.
M  X 2  MX2  X  F, Cl, Br, l
(iii) Reactivity towards hydrogen: All the elements except beryllium combine with hydrogen
upon heating to form their hydrides, MH2.
BeH2, however, can be prepared by the reaction of BeCl2 with LiAlH4.
2BeCl2 + LiAlH4  2BeH2 + LiCl + AlCl3
(iv) Reactivity towards acids: The alkaline earth metals readily react with acids liberating
dihydrogen.
M + 2HCl  MCl 2 + H2

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 20 of 74

(iv) Reducing nature: Like alkali metals, the alkaline earth metals are strong reducing agents.
This is indicated by large negative values of their reduction p ote nti als (Table 10.2).
However, their reducing power is less than those of their corresponding alkali metals. Beryllium
has less negative value compared to other alkaline earth metals. However, its reducing nature
is due to large hydration energy associated with the small size of Be2+ ion and relatively large
value of the atomization enthalpy of the metal.
(v) Solutions in liquid ammonia: Like alkali metals, the alkaline earth metals dissolve i n
liquid ammonia to give deep blue black solutions forming ammoniated ions.

M   x  y NH3  M(NH3)x]2+  2 [e(NH3)y]-

From these solutions, the ammoniates, M(NH3)x]2+ can be recovered.

10.4 USES

Beryllium is used in the manufacture of alloys. Copper-beryllium alloys are used in the preparation of high
strength springs. Metallic beryllium is used for making windows of X-ray tubes. Magnesium forms alloys
with aluminium, zinc, manganese and tin. Magnesium-aluminium alloys being light in mass are used in air-
craft construction. Magnesium (powder and ribbon) is used in flash powders and bulbs, incendiary bombs
and signals. A suspension of magnesium hydroxide in water (called milk of magnesia) is used as antacid in
medicine. Magnesium carbonate is an ingredient of toothpaste. Calcium is used in the extraction of metals
from oxides which are difficult to reduce with carbon. Calcium and barium metals, owing to their reactivity
with oxygen and nitrogen at elevated temperatures, have often been used to remove air from vacuum tubes.
Radium salts are used in radiotherapy, for example, in the treatment of cancer.

10.5 ANOMOLOUS BEHAVIOUR OF BERYLLIUM

Beryllium, the first member of the Group 2 metals, shows anomalous behavior as compared to magnesium
and rest of the members.

(i) Beryllium has exceptionally small atomic and ionic sizes and thus not well with other
members of group. Because of high ionisation enthalpy and small size it forms compounds
which are largely covalent and get easily hydrolysed.
(ii) Beryllium does not exhibit coordination number more than four as in its valence shell there
are only four orbitals. The Remaining members of the group can have a coordination
number of six by making use of d-orbitals.
(iii) The oxide and hydroxide of beryllium, unlike the hydroxides other elements in the group,
are amphoteric in nature.

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 21 of 74

. co s e m
m la

se ag

Adapted from NCERT textbook of class XII 2006 edition reprinted edition pre rationalization

m
m .co
m ELEMENTS .co s e m
se
THE p-BLOCK g l a
g la a
a
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
2 1-6
groups. Their valence shell electronic configuration is ns np (except for He). The inner core of
the electronic configuration may, however, differ. The difference in inner core of elements

m
greatly influences their physical properties (such as atomic and ionic radii, ionisation enthalpy, etc.)

.co
as well as chemical properties. Consequently, a lot of variation in properties of elements in a group

m
of p-block is observed. The maximum oxidation state shown by a p-block element is equal to the
e
l as
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
ag
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 is sometime attributed to the
‘inert pair effect’.

m
m
c. o Group
Table 11.1 General Electronic Configuration and Oxidation States of p-Block Elements

m.co
m General
13 14 15 16
s
17
e 18

s e l a
l a ns2np1 ns2np2 ns2np3 ns2np4
agns np
2 5
ns2np6

ag
electronic
configuration (1s2 for He)

First member
of the B C N O F He
group
Group
oxidation +3 +4 +5 +6 +7 +8
state

m
m .co s e m

se
For more Question Papers, Sample Papers, Notes & Syllabus visit la
Page 22 of 74

Other
oxidation +1 +2, – 4 +3, – 3 +4, +2, –2 +5, + 3, +1, –1 +6, +4, +2
states

The relative stabilities of these two oxidation states – group oxidation state and two units 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 f o r m e d
b e t w e e n n o n -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-block3
differs from the remaining members of their corresponding group in two
major respects.
4
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 3s23p 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
[BF4]-, aluminium gives [AlF6]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  bonds. The first member
of a group differs from the heavier members in its ability to form p - p multiple bonds to itself ( e.g.,
C=C, CC NN) and to other second row elements (e.g., C=O, C=N, CN, N=O). This type of  - bonding
is not particularly strong for the heavier p-block elements. The heavier elements do form  bonds
but this involves d orbitals (d – p or d –d ). As the d orbitals are of higher energy than the p
orbitals, they contribute less to the overall stability of molecules than does p - p 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 : NO3 (three-coordination with  – bond involving one nitrogen p-orbital)
and PO4 (four-coordination involving s, p and d orbitals contributing to the  – bond). In this
3

unit we will study the chemistry of group 13 and 14 elements of the periodic table.

GROUP 13 ELEMENTS: THE BORON FAMILY
This group element shows 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 BO3 ), borax,

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 23 of 74

Na2B4O7·10H2O, and kernite, Na2B4O7·4H2O. 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 10B (19%) and 11B (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, Al2 O3. 2H2 O and cryolite, Na 3 AlF 6 are the important
m i n e r a l s 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] 5f14
6d10 7s2 7p2. So far it has been prepared 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.
Electronic Configuration
2 1
The outer electronic configuration of these elements is ns np . 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.
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 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 smoothly
down the group. The decrease from B to Al is associated with increased 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 iH1 <iH2<iH3. The sum of the
in small amount and half-life of its most stable first three ionisation enthalpies for each of the
elements is very high. Effect of this will be apparent when you study their chemical properties

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.
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.

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 24 of 74

. co s e m
m la

se ag

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
m
.co
Al, the sum of the first three ionisation enthalpies of Al considerably decreases, and is

m
therefore able to form Al3+ ions. In fact, aluminium is a highly electropositive metal. However,

.co e m
down the group, due to poor shielding effect of intervening d and f orbitals, the increased

e m
effective nuclear charge holds ns electrons tightly (responsible for inert pair effect) and
l as
l as
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
ag
ag
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
Element
Property Boron Aluminium Gallium Indium Thallium
B Al Ga In Tl
Atomic number 5 13 31 49 81
–1
Atomic mass(g mol ) 10.81 26.98 69.72 114.82 204.38
Electronic [He]2s22p1 [Ne]3s23p1
o m
[Ar]3d104s24p1 [Kr]4d105s25p1 [Xe]4f145d106s26p1

135 .c
Configuration

m
a
Atomic radius/pm (88) 143 167 170
Ionic radius (27) 53.5
s e62.0 80.0 88.5
M3+/pmb
g la 120
Ionic radius
M+/pm
- -
a 140 150

Ionization iH 1 801 577 579 558 589
enthalpy iH 2 2427 1816 1979 1820 1971
(kJ mol–1) H 3659 2744 2962 2704 2877
i 3

Electronegativityc 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

m
co
Boiling point / K 3923 2740 2676 2353 1730

m
EV/ V for (M3+/M) - –1.66 –0.56 –0.34
.
+1.26

.co EV / V for (M+/M) -

m Metallic radius, 6-coordination, Pauling scale,
+0.55 -0.79(acid)
s
–1.39(alkali)
–0.18
em
–0.34

s e g l a
a
b c

a
a

l
ag 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 BF3) 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. BCl3 easily accepts a
lone pair of electrons from ammonia to form BCl3NH3
m
m .co s e m

se
For more Question Papers, Sample Papers, Notes & Syllabus visit la
Page 25 of 74

.

AlCl3 achieves stability by forming a dimer

In trivalent state most of the compounds being covalent are hydrolysed in water. For example,
-
the trichlorides on hyrolysis in water form tetrahedral [M(OH)4] species; the hybridisation state of
element M is sp3. Aluminium chloride in acidified aqueous solution forms octahedral [Al(H2O)6]3+
ion.


(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 B2O3 and Al2 O3 respectively. With dinitrogen at high temperature they
form nitrides.

2E  s   3O2  g    2E2O 3 s
2E  s   N2  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.
3+ –
2Al(s) + 6HCl (aq)  2Al (aq) + 6Cl (aq)+ 3H2(g)

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 26 of 74

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) + 6H2O(l)  2 Na [Al(OH)4] (aq) + 3H2 (g)
Sodium tetrahydroxoaluminate(III)
(iii) Reactivity towards halogens
These elements react with halogens to form trihalides (except TlI 3).
2E(s) + 3 X2 (g)  2EX3 (s) (X = F, Cl, Br, I)

IMPORTANT TRENDS AND ANOMALOUS PROPERTIES OF
BORON
Certain important trends can be observed in t h e c h e m i c a l b e h a v i o u r o f g r o u p 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(H2O)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 NH3 to complete octet around boron.
F3B :NH3 F3 B  NH3
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., AlCl3) are dimerised through halogen bridging (e.g.,
Al2Cl6). The metal species completes its octet by accepting electrons from halogen in these
halogen bridged molecules.

GROUP 14 ELEMTS: THE CABON FAMILY
Carbon, silicon, germanium, tin led and flerovium are the member 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

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 27 of 74

o m
morganisms. Naturally occurring carbon contains two stable isotopes: C and c
. C. In addition to these,

c. o third s
m
e 5770 years and used for
12 13

m a
se l
14
isotope, C is also present. It is a radioactive isotope with half-life
radiocarbon dating. Silicon is the second (27.7% by mass) mostgabundant
a is a veryelement
crust and is present in nature in the form of silica and silicates. Silicon
on the earth’s
important component
of ceramics, glass and cement. Germanium exists only in traces. Tin occurs mainly as cassiterite,
SnO2 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 146d107s2 7p2. 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
m
.co
properties along with their electronic configuration of the elements of group 14 leaving
m
.co
flerovium are given in Table 11.3. Some of the atomic, physical and chemical properties are
discussed below:
e m
e m l as
l s
Electronic Configuration
The valenceashell electronic configuration of these elements is ns np . The inner core of
a gthe 2 2

ag configuration of elements in this group also differs.
electronic
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.

Table 11.3 Atomic and Physical Properties of Group 14 Elements
Element

m
Property Carbon Silicon Germanium Tin Lead

.co
C Si Ge Sn Pb
Atomic Number 6 14 32 50 82

em
–1

s
Atomic mass (g mol ) 12.01 28.09 72.60 118.71 207.2

la
2 2 2 2 10 2 2 10 2 2 14 10 2 2
Electronic configuration [He]2s 2p [Ne]3s 3p [Ar]3d 4s 4p [Kr]4d 5s 5p [Xe]4f 5d 6s 6p

Covalent radius/pm
a
77a
g 118 122 140 146
–
4+ b
Ionic radius M /pm 40 53 69 78
– –
2+ b
Ionic radius M /pm 73 118 119
Ionization H 1086 786 761 708 715
i 1
enthalpy/ H 2352 1577 1537 1411 1450
i 2
kJ mol–1 H 4620 3228 3300 2942 3081
i 3
H 6220 4354 4409 3929 4082
i 4

om
c
Electronegativity 2.5 1.8 1.8 1.8 1.9

c
d –3 e f

m .
Density /g cm 3.51 2.34 5.32 7.26 11.34

c. o
Melting point/K 4373

s e
1693 1218
m505 600

m Electrical resistivity/ohm cm 10 –10
–
a
Boiling point/K 3550 3123 2896 2024

s e 14

g
16
l –5 –5

la
50 50 10 2 × 10

g
(293 K)
a
a a IV b c d
for M oxidation state; 6–coordination; Pauling scale; 293 K; for diamond; for graphite, density
is
e

f
2.22; -form (stable at room temperature)

m .
.co s e m
s em l a
g la ag
a For more Question Papers, Sample Papers, Notes & Syllabus visit Page 28 of 74

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 iH from Si to Ge to Sn and
slight increase in iH from Sn to Pb is the consequence of poor shielding effect of
intervening d and f orbitals and increase in size of the atom.

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.

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.

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 ns2 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 CCl4) 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, SiF6 2–, [GeCl6]2–,
[Sn(OH)6]2- exist where the hybridization of central atom is sp3d2.

(i) Reactivity towards o x y g e n

All members when heated in oxygen form oxides. There are mainly two types of oxides, i.e.,
monoxide and dioxide of the formula MO and MO2 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
— CO2, SiO2 and GeO2 are acidic, whereas SnO2 and PbO2 are amphoteric in nature.
Among monoxides, CO is neutral, GeO is distinctly acidic whereas SnO and PbO are
amphoteric.

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 29 of 74

(ii) Reactivity towards water
Carbon, silicon and germanium are not affected by water. Tin decomposes steam to form
dioxide and dihydrogen gas.

Sn + 2H2 O SnO2 + 2H2
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 MX4 are covalent in nature. The central metal atom in these
halides undergoes sp3 hybridisation and the molecule is tetrahedral in shape. Exceptions
are SnF4 and PbF4, 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 6s2 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 MX2. Stability of dihalides increases
down the group.

Considering the thermal and chemical stability, GeX4 is more stable than GeX2 , whereas PbX2 is more
than PbX4 . Except CCl4 , other tetrachlorides are easily hydrolysed by water because the central atom
can accommodatethe 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 :

+

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. Carbon
also has unique ability to form p– p multiple bonds with itself and with other atoms of small

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 30 of 74

o m
m and high electronegativity. Few examples of multiple bonding are:mC=C, c
. C  C, C = O, C

c. o size s e atomic orbitals are too
m a
se
= S, and C  N. Heavier elements do not form p– p bonds because their
large and diffuse to have effective overlapping.
g l
a
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.
Due to property of catenation and p– p bond formation, carbon is able to show allotropic
forms.
m
m .co
m .co s e m
s e l a
g l a ag
a

Adapted from NCERT textbook of class XII 2006 edition reprinted edition pre rationalization
m
Group 15 Elements
m .co
s e
la
Group 15 includes nitrogen, phosphorus, arsenic, antimony and bismuth. As we

ag
go down the group, there is a shift from non-metallic to metallic through
metalloidic character. Nitrogen and phosphorus are non-metals, arsenic and
antimony metalloids and bismuth is a typical metal.

Occurrence

Molecular nitrogen comprises 78% by volume of the atmosphere. In the earth’s
crust, it occurs as sodium nitrate, NaNO 3 (called Chile saltpetre) and potassium
m
.co
nitrate (Indian saltpetre). It is found in the form of proteins in plants and animals.

m
Phosphorus occurs in minerals of the apatite family, Ca9(PO4)6. CaX2 (X = F, Cl or

.co e m
OH) (e.g., fluorapatite Ca9 (PO4)6. CaF2) which are the main components of

e m
phosphate rocks. Phosphorus is an essential constituent of animal and plant
l as
las matter. It is present in bones as well as in living cells. Phosphoproteins are present
in milk and eggs. Arsenic, antimony and bismuth are found mainly as sulphide
ag
ag minerals.

m .
.co s e m
s em l a
g la ag
a For more Question Papers, Sample Papers, Notes & Syllabus visit Page 31 of 74

The important atomic and physical properties of this group elements
along with their electronic configurations are given in Table 7.1.

Table 7.1: Atomic and Physical Properties of Group 15 Elements

a EIII single bond (E = element); b E 3–; c E 3+; d White phosphorus; e Grey -form at 38.6 atm; f
Sublimation temperature;
g At 63 K; h Grey -form; * Molecular N .
2

Trends of some of the atomic, physical and chemical properties of the
group are discussed below.

Electronic Configuration The valence shell electronic configuration of these elements is ns2np3.
The s orbital in these elements is completely filled and p orbitals are
half-filled, making their electronic configuration extra stable.

Atomic and Ionic Radii Covalent and ionic (in a particular state) radii increase in size down
the group. There is a considerable increase in covalent radius from N
to P. However, from As to Bi only a small increase in covalent radius
is observed. This is due to the presence of completely filled d
and/or f orbitals in heavier members.

Ionisation Enthalpy Ionisation enthalpy decreases down the group due to gradual increase
in atomic size. Because of the extra stable half-filled p orbitals electronic
configuration and smaller size, the ionisation enthalpy of the group 15
elements is much greater than that of group 14 elements in the
corresponding periods. The order of successive ionisation enthalpies,
as expected is iH1 < iH2 < iH3 (Table 7.1).

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 32 of 74

Electronegativity The electronegativity value, in general, decreases down the group with
increasing atomic size. However, amongst the heavier elements, the
difference is not that much pronounced.
Physical Properties All the elements of this group are polyatomic. Dinitrogen is a diatomic gas
while all others are solids. Metallic character increases down the group.
Nitrogen and phosphorus are non-metals, arsenic and antimony metalloids
and bismuth is a metal. This is due to decrease in ionisation enthalpy and
increase in atomic size. The boiling points, in general, increase from top to
bottom in the group but the melting point increases upto arsenic and then
decreases upto bismuth. Except nitrogen, all the elements show allotropy.

Chemical Properties Oxidation states and trends in chemical reactivity
The common oxidation states of these elements are –3, +3 and +5. The
tendency to exhibit –3 oxidation state decreases down the group due to
increase in size and metallic character. In fact, last member of the group,
bismuth hardly forms any compound in –3 oxidation state. The stability
of +5 oxidation state decreases down the group. The only well characterized
Bi (V) compound is BiF5. The stability of +5 oxidation state decreases and
that of +3 state increases (due to inert pair effect) down the group. Nitrogen
exhibits + 1, + 2, + 4 oxidation states also when it reacts with oxygen.
Phosphorus also shows +1 and +4 oxidation states in some oxoacids.
In the case of nitrogen, all oxidation states from +1 to +4 tend to
disproportionate in acid solution. For example,
3HNO2  HNO3 + H2 O + 2NO
Similarly, in case of phosphorus nearly all intermediate oxidation
states disproportionate into +5 and –3 both in alkali and acid. However
+3 oxidation state in case of arsenic, antimony and bismuth becomes
increasingly stable with respect to disproportionation.
Nitrogen is restricted to a maximum covalency of 4 since only four
(one s and three p) orbitals are available for bonding. The heavier elements
have vacant d orbitals in the outermost shell which can be used for
bonding (covalency) and hence, expand their covalence as in PF6–.
Anomalous properties of nitrogen
Nitrogen differs from the rest of the members of this group due to its
small size, high electronegativity, high ionisation enthalpy and non-
availability of d orbitals. Nitrogen has unique ability to form p -p
multiple bonds with itself and with other elements having small size
and high electronegativity (e.g., C, O). Heavier elements of this group
do not form p -p bonds as their atomic orbitals are so large and
diffuse that they cannot have effective overlapping. Thus, nitrogen
exists as a diatomic molecule with a triple bond (one s and two p)
between the two atoms. Consequently, its bond enthalpy (941.4 kJ mol–
1
) is very high. On the contrary, phosphorus, arsenic and antimony
form single bonds as P–P, As–As and Sb–Sb while bismuth forms
metallic bonds in elemental state. However, the single N–N bond is weaker
than the single P–P bond because of high interelectronic repulsion of the
non-bonding electrons, owing to the small bond length. As a result, the
catenation tendency is weaker in nitrogen.

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 33 of 74

m
m .co

m .co s em
se a nitrogen is the absence of d
Another factor which affects the chemistrylof
orbitals in its valence shell. Besides g
a restricting its covalency to four,
nitrogen cannot form d –p bond as the heavier elements can e.g., R P 3
= O or R3P = CH2 (R = alkyl group). Phosphorus and arsenic can form d –
d bond also with transition metals when their compounds like P(C2H5)3
and As(C6 H5)3 act as ligands.
(i) Reactivity towards hydrogen: All the elements of Group 15 form
hydrides of the type EH3 where E = N, P, As, Sb or Bi. Some of
the properties of these hydrides are shown in Table 7.2. The
hydrides show regular gradation in their properties. The
m
.co
stability of hydrides decreases from NH3 to BiH3 which can be
m
.co
observed from their bond dissociation enthalpy.
Consequently, the reducing character of the hydrides increases.
e m
em Ammonia is only a mild reducing agent while BiH3 is the
l as
l as strongest reducing agent amongst all the hydrides. Basicity also
ag
g
decreases in the order NH3 > PH3 > AsH3 > SbH3 > BiH3.
a Table 7.2: Properties of Hydrides of Group 15 Elements

Melting point/K 195.2 139.5 156.7 185 –

Boiling point/K 238.5 185.5 210.6 254.6 290

(E–H) Distance/pm 101.7 141.9 151.9 170.7 –

–
m
HEH angle (°) 107.8 93.6 91.8 91.3

 H / kJ mol
-1

c. o 13.4 66.4 145.1 278
–46.1

 H (E–H)/kJ mol em
f

s –
–1
389 322 297 255

la oxygen: All these elements form two types
diss

a
(ii) Reactivity gtowards
of oxides: E O and E O . The oxide in the higher oxidation state of
2 3 2 5
the element is more acidic than that of lower oxidation state. Their
acidic character decreases down the group. The oxides of the type
E2O3 of nitrogen and phosphorus are purely acidic, that of
arsenic and antimony amphoteric and those of bismuth
predominantly basic.
(iii) Reactivity towards halogens: These elements react to form two
series of halides: EX3 and EX5. Nitrogen does not form
m
.co
pentahalide due to non-availability of the d orbitals in its valence
m
.co
shell. Pentahalides are more covalent than trihalides. All the

e m
trihalides of these elements except those of nitrogen are stable. In

e m as
case of nitrogen, only NF3 is known to be stable. Trihalides except
l
las BiF3 are predominantly covalent in nature.
ag
g
(iv) Reactivity towards metals: All these elements react with metals to
a form their binary compounds exhibiting –3 oxidation state, such
as, Ca3N2 (calcium nitride) Ca3P2 (calcium phosphide), Na3As2
(sodium arsenide), Zn3Sb2 (zinc antimonide) and Mg3Bi2
(magnesium bismuthide).

m .
.co s e m
s em l a
g la ag
a For more Question Papers, Sample Papers, Notes & Syllabus visit Page 34 of 74

Though nitrogen exhibits +5 oxidation state, it does not form Example 7.1
pentahalide. Give reason.
Nitrogen with n = 2, has s and p orbitals only. It does not have d orbitals Solution
to expand its covalence beyond four. That is why it does not form
pentahalide.

PH3 has lower boiling point than NH3. Why? Example 7.2
Unlike NH3, PH3 molecules are not associated through hydrogen bonding Solution
in liquid state. That is why the boiling point of PH3 is lower than NH3.

Questions

Group 16 Elements

Oxygen, Sulphur, selenium, tellurium and polonium constitute Group 16 of the periodic table.
This is sometimes known as group of chalcogens. The name is derived from the Greek word
for brass and points to the association of Sulphur and its congeners with copper.
Most copper minerals contain either oxygen or Sulphur and frequently the other members
of group.

Occurrence
Oxygen is the most abundant of the elements on earth. Oxygen forms about
46.6%by mass earth’s crust. Dry air contains 20.946% oxygen by volume. However,
the abundance of Sulphur in the earth’s crust I only 0.03-0.1%. Combined Sulphur
exists primarily as sulphates such as gypsum CaSO4.2H2O, Epsom salt MgSO4.7H2O,
barite BaSO4 and sulphides such as galena PbS, zinc blende ZnS, copper pyrites
CuFeS2. Traces of Sulphur occur as hydrogen sulphide in volcanoes, Organic
materials such as eggs, proteins, garlic, onion, mustard, hair and wool contain
Sulphur Selenium and tellurium are also found as metal selenides and
tellurides in sulphide ores. Polonium occurs in nature as a decay
product of thorium and uranium minerals. The important atomic and
physical properties of Group16 along with electronic configuration are
given in Table 7.3. Some of the atomic, physical and chemical
properties and their trends are discussed below.

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 35 of 74

Electronic The elements of Group16 have six electrons in the outermost shell and
Configuration have ns2np4 general electronic configuration.

Atomic and Due to increase in the number of shells, atomic and ionic radii increase
Ionic Radii from top to bottom in the group. The size of oxygen atom is, however,
exceptionally small.

Ionisation enthalpy decreases down the group. It is due to increase in
Ionisation size. However, the elements of this group have lower ionisation enthalpy
Enthalpy values compared to those of Group15 in the corresponding periods.
This is due to the fact that Group 15 elements have extra stable half-
filled p orbitals electronic configurations.
Electron Gain
Enthalpy Because of the compact nature of oxygen atom, it has less negative
electron gain enthalpy than sulphur. However, from sulphur onwards
the value again becomes less negative upto polonium.

Electronegativity Next to fluorine, oxygen has the highest electronegativity value amongst
the elements. Within the group, electronegativity decreases with an
increase in atomic number. This implies that the metallic character
increases from oxygen to polonium.

Elements of Group 16 generally show lower value of first ionisation Example 7.3
enthalpy compared to the corresponding periods of group 15. Why?
Due to extra stable half-filled p orbitals electronic configurations of Solution
Group 15 elements, larger amount of energy is required to remove
electrons compared to Group 16 elements.

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 36 of 74

m
m .co

m .co Physical Properties
s e m
se l a
Some of the physical properties of Group 16 elements are given in Table

ag is radioactive and is
7.3 Oxygen and sulphur are non-metals, selenium
metalloids, whereas polonium is a metal. Polonium
and tellurium

short lived (Half-life 13.8 days). All these elements exhibit allotropy. The
melting and boiling points increase with an increase in atomic number
down the group. The large difference between the melting and boiling
points of oxygen and sulphur may be explained on the basis of their
atomicity; oxygen exists as diatomic molecule (O 2) whereas sulphur
exists as polyatomic molecule (S 8).

m
.co
Chemical Properties

m
.co m
Oxidation states and trends in chemical reactivity

m
The elements of Group 16 exhibit a number of oxidation states (Table 7.3). The
s e
s e
stability of -2 oxidation state decreases down the group. Polonium hardly shows
l a
g l a
–2 oxidation state. Since electronegativity of oxygen is very high, it shows only
negative oxidation state as –2 except in the case of OF2 where its oxidation ag
a state is + 2.

Table 7.3: Some Physical Properties of Group 16 Elements

m
m .co
s e
g la
a

m
m .co
m .co s e m
s e g l a
g la a
a

m .
.co s e m
s em l a
g la ag
a For more Question Papers, Sample Papers, Notes & Syllabus visit Page 37 of 74

Other elements of the group exhibit + 2, + 4, + 6 oxidation states but + 4 and + 6 are more
common. Sulphur, selenium and tellurium usually show + 4 oxidation state in their
compounds with oxygen and + 6 with fluorine. The stability of + 6 oxidation state decreases
down the group and stability of + 4 oxidation state increases (inert pair effect). Bonding in +4
and +6 oxidation states is primarily covalent.
Anomalous behaviour of oxygen
The anomalous behaviour of oxygen, like other members of p-block present in second period
is due to its small size and high electronegativity. One typical example of effects of small size
and high electronegativity is the presence of strong hydrogen bonding in H 2 O which is not
found in H2S.
The absence of d orbitals in oxygen limits its covalency to four and in practice, rarely exceeds
two. On the other hand, in case of other elements of the group, the valence shells can be
expanded and covalence exceeds four.
(i) Reactivity with hydrogen: All the elements of Group 16 form hydrides of the type
H2E (E = O, S, Se, Te, Po). Some properties of hydrides are given in Table 7.4.
Their acidic character increases from H2O to H2Te. The increase in acidic character
can be explained in terms of decrease in bond enthalpy for the dissociation of H–
E bond down the group. Owing to the decrease in enthalpy for the dissociation of
H–E bond down the group, the thermal stability of hydrides also decreases from
H2O to H2Po. All the hydrides except water possess reducing property and this
character increases from H2S to H2Te.
(ii) Reactivity with oxygen: All these elements form oxides of the EO 2 and EO3 types
where E = S, Se, Te or Po. Ozone (O 3) and sulphur dioxide (SO2) are gases while
selenium dioxide (SeO2) is solid. Reducing property of dioxide decreases from SO 2
to TeO2; SO2 is reducing while TeO2 is an oxidising agent. Besides EO2 type,
sulphur, selenium and tellurium also form EO3 type oxides (SO3, SeO3, TeO3).
Both types of oxides are acidic in nature.
(iii) Reactivity towards the halogens: Elements of Group 16 form a large number of
halides of the type, EX6, EX4 and EX2 where E is an element of the group and X is a
halogen. The stability of the halides decreases in the order F – > Cl – > Br– > I–.
Amongst hexahalides, hexafluorides are the only stable halides. All hexafluorides
are gaseous in nature. They have octahedral structure. Sulphur hexafluoride, SF 6
is exceptionally stable for steric reasons.
Amongst tetrafluorides, SF4 is a gas, SeF4 a liquid and TeF4 a solid. These fluorides have
sp3d hybridisation and thus, have trigonal bipyramidal structures in which one of the
equatorial positions is occupied by a lone pair of electrons. This geometry is also regarded
as see-saw geometry.
All elements except oxygen form dichlorides and dibromides. These dihalides are formed
by sp3 hybridisation and thus, have tetrahedral structure. The well-known monohalides are
dimeric in nature. Examples are S2F2, S2Cl2, S2Br2, Se2Cl2 and Se2Br2. These dimeric halides
undergo disproportionation as given below:
2Se2 Cl2  SeCl4 + 3Se

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 38 of 74

Table 7.4: Properties of Hydrides of Group 16 Elements

Property H2O H2S H 2Se H2Te

m.p/K 273 188 208 222
b.p/K 373 213 232 269
H–E distance/pm 96 134 146 169
HEH angle (°) 104 92 91 90

a Aqueous solution, 298 K

H2S is less acidic than H2Te. Why? Example7.4
Due to the decrease in bond (E–H) dissociation Solution
enthalpy down the group, acidic character increases.

Questions

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 39 of 74

m
m .co

m .co s e m
se g l a
a
Group 17 Fluorine, c h l o r i n e , b r o m i n e , i o d i n e a n d a s t a t i n e a r e
members of
Elements Group 17. These are collectively known as the halogens (Greek

halo means salt and genes means born i.e., salt producers). The
halogens are highly reactive non-metallic elements. Like Groups 1 and
2, the elements of Group 17 show great similarity amongst themselves.
That much similarity is not found in the elements of other groups
m
m .co
of the periodic table. Also, there is a regular gradation in their physical

. c o and chemical properties. Astatine is a radioactive element.
e m
Occurrence m Fluorine and chlorine are fairly abundant while bromine and iodine less
s e so. Fluorine is present mainly as insoluble fluorides (fluorspargCaF l as,
la a
2

a g cryolite Na AlF and fluoroapatite 3Ca (PO ) .CaF and small quantities
3 6 3 4 2
are present in soil, river water plants and bones and teeth of animals.
2

Sea water contains chlorides, bromides and iodides of sodium,
potassium, magnesium and calcium, but is mainly sodium chloride
solution (2.5% by mass). The deposits of dried up seas contain these
compounds, e.g., sodium chloride and carnallite, KCl.MgCl 2.6H2 O.
Certain forms of marine life contain iodine in their systems; various
seaweeds, for example, contain upto 0.5% of iodine and Chile saltpetre
contains upto 0.2% of sodium iodate.
The important atomic and physical properties of Group 17 elements
m
.co
along with their electronic configurations are given in Table 7.5.

e m
s
Table 7.5: Atomic and Physical Properties of Halogens

g la
a

m
m .co
.co
–

m
F2 Cl2 Br2 I2

s e
em
Melting point/K 54.4 172.0 265.8 386.6 –

s
Boiling point/K 84.9 239.0 332.5

g l a
458.2 –

l a a –

ag
Distance X – X/pm 143 199 228 266 –
Bond dissociation enthalpy 158.8 242.6 192.8 151.1 –
/(kJ mol–1)
EV/Ve 2.87 1.36 1.09 0.54 –

a Radioactive; b Pauling scale; c For the liquid at temperatures (K) given in the parentheses; d solid; e The
half-cell reaction is X (g)
2 + 2e–  2X–(aq).

m .
.co s e m
s em l a
g la ag
a For more Question Papers, Sample Papers, Notes & Syllabus visit Page 40 of 74

The trends of some of the atomic, physical and chemical properties are
discussed below.

Electronic All these elements have seven electrons in their outermost shell
Configuration (ns2np5) which is one electron short of the next noble gas.

Atomic The halogens have the smallest atomic radii in their respective periods due
and Ionic to maximum effective nuclear charge. The atomic radius of fluorine like
Radii the other elements of second period is extremely small. Atomic and ionic
radii increase from fluorine to iodine due to increasing number of
quantum shells.

Ionisation They have little tendency to lose electron. Thus they have very high
Enthalpy ionisation enthalpy. Due to increase in atomic size, ionisation enthalpy
decreases down the group.

Electron Halogens have maximum negative electron gain enthalpy in the
Gain corresponding periods. This is due to the fact that the atoms of these
Enthalpy elements have only one electron less than stable noble gas configurations.
Electron gain enthalpy of the elements of the group becomes less negative
down the group. However, the negative electron gain enthalpy of fluorine
is less than that of chlorine. It is due to small size of fluorine atom. As
a result, there are strong interelectronic repulsions in the relatively
small 2p orbitals of fluorine and thus, the incoming electron does not
experience much attraction.

Electronegativity They have very high electronegativity. The electronegativity decreases
down the group. Fluorine is the most electronegative element in the
periodic table.

Halogens have maximum negative electron gain enthalpy in the Example 7.5
respective periods of the periodic table. Why?
Halogens have the smallest size in their respective periods and therefore Solution
high effective nuclear charge. As a consequence, they readily accept
one electron to acquire noble gas electronic configuration.

Physical Halogens display smooth variations in their physical properties. Fluorine
Properties and chlorine are gases, bromine is a liquid and iodine is a solid. Their
melting and boiling points steadily increase with atomic number. All
halogens are coloured. This is due to absorption of radiations in visible
region which results in the excitation of outer electrons to higher energy
level. By absorbing different quanta of radiation, they display different
colours. For example, F2, has yellow, Cl2 , greenish yellow, Br2, red and
I2, violet colour. Fluorine and chlorine react with water. Bromine and
iodine are only sparingly soluble in water but are soluble in various
organic solvents such as chloroform, carbon tetrachloride, carbon
disulphide and hydrocarbons to give coloured solutions.

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 41 of 74

One curious anomaly we notice from Table 7.5 is the smaller enthalpy of
dissociation of F2 compared to that of Cl2 whereas X-X bond dissociation enthalpies
from chlorine onwards show the expected trend: Cl – Cl > Br – Br > I – I. A reason for
this anomaly is the relatively large electron-electron repulsion among the lone pairs
in F2 molecule where they are much closer to each other than in case of Cl2.

Example 7.6 Although electron gain enthalpy of fluorine is less negative as compared
to chlorine, fluorine is a stronger oxidising agent than chlorine. Why?
Solution It is due to
(i) low enthalpy of dissociation of F-F bond (Table 7.8).
(ii) high hydration enthalpy of F– (Table 7.8).

Chemical Oxidation states and trends in chemical reactivity
Properties All the halogens exhibit –1 oxidation state. However, chlorine, bromine
and iodine exhibit + 1, + 3, + 5 and + 7 oxidation states also as explained
below:

Halogen atom ns np nd
in ground state 1 unpaired electron accounts
(other than fluorine) for –1 or +1 oxidation states

3 unpaired electrons account
First excited state
for +3 oxidation states
5 unpaired electrons account
Second excited state
for +5 oxidation state
Third excited state 7 unpaired electrons
account for +7 oxidation state

The higher oxidation states of chlorine, bromine and iodine are realised
mainly when the halogens are in combination with the small and highly
electronegative fluorine and oxygen atoms. e.g., in interhalogens, oxides
and oxoacids. The oxidation states of +4 and +6 occur in the oxides and
oxoacids of chlorine and bromine. The fluorine atom has no d orbitals
in its valence shell and therefore cannot expand its octet. Being the most
electronegative, it exhibits only –1 oxidation state.
All the halogens are highly reactive. They react with metals and
non-metals to form halides. The reactivity of the halogens decreases down
the group.
The ready acceptance of an electron is the reason for the strong
oxidising nature of halogens. F2 is the strongest oxidising halogen and
it oxidises other halide ions in solution or even in the solid phase. In
general, a halogen oxidises halide ions of higher atomic number.
F2 + 2X–  2F– + X2 (X = Cl, Br or I)
Cl2 + 2X–  2Cl– + X2 (X = Br or I)
Br2 + 2I–  2Br– + I2
The decreasing oxidising ability of the halogens in aqueous solution
down the group is evident from their standard electrode potentials (Table
7.5) which are dependent on the parameters indicated below:

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 42 of 74

m
m .co

m .co s e m
se g l a
a
1/ 2 diss HV egH V hydH V –
–
1 / 2 X 2 g  X g  X  X aq
2

The relative oxidising power of halogens can further be illustrated
by their reactions with water. Fluorine oxidises water to oxygen whereas
chlorine and bromine react with water to form corresponding hydrohalic
and hypohalous acids. The reaction of iodine with water is non-
spontaneous. In fact, I – can be oxidised by oxygen in acidic medium;
m
m .co
just the reverse of the reaction observed with fluorine.

c. o 2F g   2H Ol  4H aq  4F aq  O g
 
s e m
em 2 2 2

l a
l as X 2  g   H 2 Ol  HX aq  HOX aq
ag
ag where X = Cl or Br
4I  aq   4H  aq  O 2  g   2I2 s  2H 2O l 

Anomalous behaviour of fluorine
Like other elements of p-block present in second period of the periodic
table, fluorine is anomalous in many properties. For example, ionisation
enthalpy, electronegativity, and electrode potentials are all higher for
fluorine than expected from the trends set by other halogens. Also, ionic

m
and covalent radii, m.p. and b.p., enthalpy of bond dissociation and electron

.co
gain enthalpy are quite lower than expected. The anomalous behaviour of

m
fluorine is due to its small size, highest electronegativity, low F-F bond

s e
dissociation enthalpy, and non-availability of d orbitals in valence shell.

la
Most of the reactions of fluorine are exothermic (due to the small and
g
a
strong bond formed by it with other elements). It forms only one oxoacid
while other halogens form a number of oxoacids. Hydrogen fluoride is a
liquid (b.p. 293 K) due to strong hydrogen bonding. Other hydrogen
halides are gases.

(i) Reactivity towards hydrogen: They all react with hydrogen to give
hydrogen halides but affinity for hydrogen decreases from fluorine
to iodine. Hydrogen halides dissolve in water to form hydrohalic
acids. Some of the properties of hydrogen halides are given in

m
Table 7.6. The acidic strength of these acids varies in the order: HF

.co
< HCl < HBr < HI. The stability of these halides decreases down the
m
.co
group due to decrease in bond (H–X) dissociation enthalpy in the
order: H–F > H–Cl > H–Br > H–I.
e m
e m s
Table 7.6: Properties of Hydrogen aHalides
l
las a g
ag

m .
.co s e m
s em l a
g la ag
a For more Question Papers, Sample Papers, Notes & Syllabus visit Page 43 of 74

(ii) Reactivity towards oxygen: Halogens form many oxides with oxygen but
most of them are unstable. Fluorine forms two oxides OF2 and O2F2.
However, only OF2 is thermally stable at 298 K. These oxides are
essentially oxygen fluorides because of the higher electronegativity of
fluorine than oxygen. Both are strong fluorinating agents. O2F2 oxidises
plutonium to PuF6 and the reaction is used in removing plutonium
as PuF6 from spent nuclear fuel.Chlorine, bromine and iodine form
oxides in which the oxidation states of these halogens range from +1 to
+7. A combination of kinetic and thermodynamic factors leads to the
generally decreasing order of stability of oxides formed by halogens, I
> Cl > Br. The higher oxides of halogens tend to be more stable than
the lower ones.Chlorine oxides, Cl2O, ClO2, Cl2O6 and Cl2O7 are highly
reactive oxidising agents and tend to explode. ClO 2 is used as a
bleaching agent for paper pulp and textiles and in water
treatment.The bromine oxides, Br2O, BrO2 , BrO3 are the least stable
halogen oxides (middle row anomally) and exist only at low
temperatures. They are very powerful oxidising agents.The iodine
oxides, I2O4 , I2O5, I2O7 are insoluble solids and decompose on
heating. I2O5 is a very good oxidising agent and is used in the
estimation of carbon monoxide.
(iii) Reactivity towards metals: Halogens react with metals to form
metal halides. For example, bromine reacts with magnesium to give
magnesium bromide.
Mg s  Br2  l   MgBr2 s
The ionic character of the halides decreases in the order MF > MCl
> MBr > MI where M is a monovalent metal. If a metal exhibits more
than one oxidation state, the halides in higher oxidation state will
be more covalent than the one in lower oxidation state. For example,
SnCl4, PbCl4 , SbCl5 and UF6 are more covalent than SnCl2, PbCl2, SbCl3
and UF4 respectively.
(iv) Reactivity of halogens towards other halogens: Halogens combine
amongst themselves to form a number of compounds known as
interhalogens of the types XX , XX 3, XX 5 and XX 7 where X is a
larger size halogen and X is smaller size halogen.

Example 7.7 Fluorine exhibits only –1 oxidation state whereas other halogens
exhibit + 1, + 3, + 5 and + 7 oxidation states also. Explain.
Solution Fluorine is the most electronegative element and cannot exhibit any positive
oxidation state. Other halogens have d orbitals and therefore, can expand
their octets and show + 1, + 3, + 5 and + 7 oxidation states also.

Questions

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 44 of 74

7.3 Group 18 Group 18 consists of six elements: helium, neon, argon, krypton, xenon
Elements and radon. All these are gases and chemically unreactive. They form
very few compounds. Because of this they are termed noble gases.

Occurrence All the noble gases except radon occur in the atmosphere. Their
atmospheric abundance in dry air is ~ 1% by volume of which argon
is the major constituent. Helium and sometimes neon are found in
minerals of radioactive origin e.g., pitchblende, monazite, cleveite. The
main commercial source of helium is natural gas. Xenon and radon are
the rarest elements of the group. Radon is obtained as a decay product
of 226Ra.

Example 7.8 Why are the elements of Group 18 known as noble gases?
Solution The elements present in Group 18 have their valence shell orbitals
completely filled and, therefore, react with a few elements only under
certain conditions. Therefore, they are now known as noble gases.

The important atomic and physical properties of the Group 18
elements along with their electronic configurations are given in
Table 7.7. The trends in some of the atomic, physical and chemical
properties of the group are discussed her.

Table 7.7:Atomic and Physical Properties of Group 18 Elements
* radioactive

–

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 45 of 74

m
m .co

m .co s em
se
Electronic
g a
All noble gases have general electronic lconfiguration 2 6

helium which has 1s (Table 7.7). Manyaof the properties of noble gases
Configuration ns np except
2

including their inactive nature are ascribed to their closed shell
structures.
Ionisation
Due to stable electronic configuration these gases exhibit very high
Enthalpy
ionisation enthalpy. However, it decreases down the group with increase
in atomic size.
Atomic
Atomic radii increase down the group with increase in atomic number.
Radii
Since noble gases have stable electronic configurations, they have no
m
.co
Electron
tendency to accept the electron and therefore, have large positive values
Gain
o m m
c
of electron gain enthalpy.
. e
Enthalpy

emAll
Physical Properties

l as
l as a g
the noble gases are monoatomic. They are colourless, odourless and

ag
tasteless. They are sparingly soluble in water. They have very low
melting and boiling points because the only type of interatomic
interaction in these elements is weak dispersion forces. Helium has the
lowest boiling point (4.2 K) of any known substance. It has an unusual
property of diffusing through most commonly used laboratory materials
such as rubber, glass or plastics.

Noble gases have very low boiling points. Why?
Noble gases being monoatomic have no interatomic forces except weak
dispersion forces and therefore, they are liquefied at very low
m
.co
temperatures. Hence, they have low boiling points.

e m
s
Chemical Properties

la
In general, noble gases are least reactive. Their inertness to chemical

ag
reactivity is attributed to the following reasons:
(i) The noble gases except helium (1s2) have completely filled
ns2np6electronic configuration in their valence shell.
(ii)They have high ionisation enthalpy and more positive electron gain
enthalpy.

The reactivity of noble gases has been investigated occasionally, ever
since their discovery, but all attempts to force them to react to form the
compounds, were unsuccessful for quite a few years. In March 1962, Neil
Bartlett, then at the University of British Columbia, observed the reaction
m
.co
of a noble gas. First, he prepared a red compound whichis formulated
m as O +PtF –. He, then realised that the first ionization enthalpy of

m .co molecular oxygen (1175 kJmol–1 ) was almost identical with that of xenon
(1170 kJmol–1). He made efforts to prepare same type of compound with s e m
s e Xe and was successful in preparing another red colour compound Xe+PtF6 –
g l a
g la by mixing PtF6 and xenon. After this discovery, a number of xenon
a
a compounds mainly with most electronegative elements like fluorine and
oxygen, have been synthesized.
The compounds of krypton are fewer. Only the difluoride (KrF2) has been
studied in detail. Compounds of radon have not been isolated but only
identified (e.g., RnF2) by radiotracer technique. No true compounds of
Ar, Ne or He are yet known.

m .
.co s e m
s em l a
g la ag
a For more Question Papers, Sample Papers, Notes & Syllabus visit Page 46 of 74

READING MATERIAL
CLASS-XII (2026-27)
CHEMISTRY (043)

Adapted from NCERT textbook of Class XII 2022-23 edition.

SURFACE CHEMISTRY

Surface chemistry deals with phenomena that occur at the surfaces or interfaces. The
interface or surface is represented by separating the bulk phases by a hyphen or a slash.
For example, the interface between a solid and a gas may be represented by solid-gas
or solid/gas. Due to complete miscibility, there is no interface between the gases. The bulk
phases that we come across in surface chemistry may be pure compounds or solutions. The
interface is normally a few molecules thick but its area depends on the size of the particles
of bulk phases. Many important phenomena, noticeable amongst these being corrosion,
electrode processes, heterogeneous catalysis, dissolution and crystallisation occur at
interfaces. The subject of surface chemistry finds many applications in industry, analytical
work and daily life situations.
To accomplish surface studies meticulously, it becomes imperative to have a really clean
surface. Under very high vacuum of the order of 10 –8 to 10–9 pascal, it is now possible to
obtain ultra clean surface of the metals. Solid materials with such clean surfaces need to
be stored in vacuum otherwise these will be covered by molecules of the major components
of air namely dioxygen and dinitrogen.
In this Unit, you will be studying some important features of surface chemistry such as
adsorption, catalysis and colloids including emulsions and gels.

5.1 Adsorption
There are several examples, which reveal that the surface of a solid has the tendency to attract
and retain the molecules of the phase with which it comes into contact. These molecules
remain only at the surface and do not go deeper into the bulk. The accumulation of
molecular species at the surface rather than in the bulk of a solid or liquid is termed
adsorption. The molecular species or substance, which concentrates or accumulates at the
surface is termed adsorbate and the material on the surface of which the adsorption takes
place is called adsorbent.
Adsorption is essentially a surface phenomenon. Solids, particularly in finely divided state,
have large surface area and therefore, charcoal, silica gel, alumina gel, clay, colloids, metals
in finely divided state, etc. act as good adsorbents.
Adsorption in action

(i) If a gas like O2, H2, CO, Cl2, NH3 or SO2 is taken in a closed vessel containing powdered
charcoal, it is observed that the pressure of the gas in the enclosed vessel decreases.
The gas molecules concentrate at the surface of the charcoal, i.e., gases are adsorbed at
the surface.
(ii) In a solution of an organic dye, say methylene blue, when animal charcoal is added
and the solution is well shaken, it is observed that the filtrate turns colourless. The
molecules of the dye, thus, accumulate on the surface of charcoal, i.e., are adsorbed.
(iii) Aqueous solution of raw sugar, when passed over beds of animal charcoal, becomes
colourless as the colouring substances are adsorbed by the charcoal.
(iv) The air becomes dry in the presence of silica gel because the water molecules get
adsorbed on the surface of the gel.
It is clear from the above examples that solid surfaces can hold the gas or liquid
molecules by virtue of adsorption. The process of removing an adsorbed substance from
a surface on which it is adsorbed is called desorption.

5.1.1 Distinction between Adsorption and Absorption
In adsorption, the substance is concentrated only at the surface and does not penetrate

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 47 of 74

through the surface to the bulk of the adsorbent, while in absorption, the substance is
uniformly distributed throughout the bulk of the solid. For example, when a chalk stick is
dipped in ink, the surface retains the colour of the ink due to adsorption of coloured molecules
while the solvent of the ink goes deeper into the stick due to absorption. On breaking the
chalk stick, it is found to be white from inside. A distinction can be made between absorption
and adsorption by taking an example of water vapour. Water vapours are absorbed by
anhydrous calcium chloride but adsorbed by silica gel. In other words, in adsorption the
concentration of the adsorbate increases only at the surface of the adsorbent, while in
absorption the concentration is uniform throughout the bulk of the solid.
Both adsorption and absorption can take place simultaneously also. The term sorption
is used to describe both the processes.

5.1.2 Types of Adsorption
There are mainly two types of adsorption of gases on solids. If accumulation of gas
on the surface of a solid occurs on account of weak van der Waals’ forces, the adsorption is
termed as physical adsorption or physisorption. When the gas molecules or atoms are
held to the solid surface by chemical bonds, the adsorption is termed chemical adsorption
or chemisorption. The chemical bonds may be covalent or ionic in nature. Chemisorption
involves a high energy of activation and is, therefore, often referred to as activated adsorption.
Sometimes these two processes occur simultaneously and it is not easy to ascertain the type
of adsorption. A physical adsorption at low temperature may pass into chemisorption as the
temperature is increased. For example, dihydrogen is first adsorbed on nickel by van der
Waals’ forces. Molecules of hydrogen then dissociate to form hydrogen atoms which are held on
the surface by chemisorption.

5.1.3 Applications of Adsorption
The phenomenon of adsorption finds a number of applications. Important ones are listed here:
(i) Production of high vacuum: The remaining traces of air can be adsorbed by charcoal
from a vessel evacuated by a vacuum pump to give a very high vacuum.
(ii) Gas masks: Gas mask (a device which consists of activated charcoal or mixture of
adsorbents) is usually used for breathing in coal mines to adsorb poisonous gases.

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 48 of 74

m
m .co

m .co s em
se g la
(iii) Control of humidity: Silica and aluminium gels are a
used as adsorbents for removing
moisture and controlling humidity.
(iv) Removal of colouring matter from solutions: Animal charcoal removes colours of
solutions by adsorbing coloured impurities.
(v) Heterogeneous catalysis: Adsorption of reactants on the solid surface of the catalysts
increases the rate of reaction. There are many gaseous reactions of industrial importance
involving solid catalysts. Manufacture of ammonia using iron as a catalyst, manufacture
of H2SO4 by contact process and use of finely divided nickel in the hydrogenation of oils
are excellent examples of heterogeneous catalysis.

m
.co
(vi) Separation of inert gases: Due to the difference in degree of adsorption of gases by charcoal,

m
a mixture of noble gases can be separated by adsorption on coconut charcoal at different

.co m
temperatures.

m
(vii) In curing diseases: A number of drugs are used to kill germs by getting adsorbed on them.
s e
s e
(viii) Froth floatation process: A low grade sulphide ore is concentrated by separating it from
l a
l a
silica and other earthy matter by this method using pine oil and frothing agent

g ag
a
(ix) Adsorption indicators: Surfaces of certain precipitates such as silver halides have the
property of adsorbing some dyes like eosin, fluorescein, etc. and thereby producing a
characteristic colour at the end point.
(x) Chromatographic analysis: Chromatographic analysis based on the phenomenon of
adsorption finds a number of applications in analytical and industrial fields.

5.2 Catalysis
Potassium chlorate, when heated strongly decomposes slowly giving dioxygen. The
decomposition occurs in the temperature range of 653-873K.

m
.co
However, when a little of manganese dioxide is added, the decomposition takes place at a
m
e
considerably lower temperature range, i.e., 473-633K and also at a much accelerated rate.

composition.
las
The added manganese dioxide remains unchanged with respect to its mass and

Promoters and poisons
ag
Promoters are substances that enhance the activity of a catalyst while poisons decrease the
activity of a catalyst. For example, in Haber’s process for manufacture of ammonia,
molybdenum acts as a promoter for iron which is used as a catalyst.

5.2.1 Homogeneous and Heterogeneous Catalysis

o m
mWhenprocess
(a) Homogeneous catalysis
c
.of liquid
c. o the m
the reactants products and the catalyst are in the same phase (i.e., or gas),

m of(i) homogeneous s e
is said to be homogeneous catalysis. The following are some the examples

s e catalysis:

g a
l in the presence of
l a a
Oxidation of sulphur dioxide into sulphur trioxide with dioxygen

ag
oxides of nitrogen as the catalyst in the lead chamber process.

The reactants, sulphur dioxide and oxygen, and the catalyst, nitric oxide, are all
in the same phase.

(b) Heterogeneous catalysis
The catalytic process in which the reactants and the catalyst are in different phases is
m .
.co
known as heterogeneous catalysis. Some of the examples of heterogeneous catalysis
are given below:
e m
em l as
las ag
ag For more Question Papers, Sample Papers, Notes & Syllabus visit Page 49 of 74

(i) Oxidation of sulphur dioxide into sulphur trioxide in the presence of Pt.

The reactant is in gaseous state while the catalyst is in the solid state.

5.2.2 Shape- Selective Catalysis by Zeolites
The catalytic reaction that depends upon the pore structure of the catalyst and the size of the
reactant and product molecules is called shape-selective catalysis. Zeolites are good
shape-selective catalysts because of their honeycomb-like structures. They are
microporous aluminosilicates with three dimensional network of silicates in which some
silicon atoms are replaced by aluminium atoms giving Al–O–Si framework. The reactions
taking place in zeolites depend upon the size and shape of reactant and product molecules
as well as upon the pores and cavities of the zeolites. They are found in nature as well as
synthesised for catalytic selectivity.
Zeolites are being very widely used as catalysts in petrochemical industries for cracking
of hydrocarbons and isomerisation. An important zeolite catalyst used in the petroleum
industry is ZSM-5. It converts alcohols directly into gasoline (petrol) by dehydrating them to
give a mixture of hydrocarbons.

5.2.3 Catalysts in Industry
Some of the important technical catalytic processes are listed in Table 5.3 to give an
idea about the utility of catalysts in industries.

5.3 Colloids
We have learnt in chapter solutions that solutions are homogeneous systems. We also know that
sand in water when stirred gives a suspension, which slowly settles down with time. Between the
two extremes of suspensions and solutions we come across a large group of systems called
colloidal dispersions or simply colloids.
A colloid is a heterogeneous system in which one substance is dispersed (dispersed phase)
as very fine particles in another substance called dispersion medium.
The essential difference between a solution and a colloid is that of particle size. While in a
solution, the constituent particles are ions or small molecules, in a colloid, the dispersed phase
may consist of particles of a single macromolecule (such as protein or synthetic polymer) or an
aggregate of many atoms, ions or molecules. Colloidal particles are larger than simple molecules
but small enough to remain suspended. Their range of diameters is between 1 and 1000 nm
(10–9 to 10–6 m).

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 50 of 74

Colloidal particles have an enormous surface area per unit mass as a result of their small
size. Consider a cube with 1 cm side. It has a total surface area of 6 cm2. If it were divided
equally into 1012 cubes, the cubes would be the size of large colloidal particles and have a total
surface area of 60,000 cm2 or 6 m2. This enormous surface area leads to some special
properties of colloids to be discussed later in this Unit.

5.4 Classification of Colloids

Colloids are classified on the basis of the following criteria:
(i) Physical state of dispersed phase and dispersion medium
(ii) Nature of interaction between dispersed phase and dispersion medium
(iii) Type of particles of the dispersed phase.

5.4.1 Classification Based on Physical State of Dispersed Phase and Dispersion
Medium

Depending upon whether the dispersed phase and the dispersion medium are solids, liquids or
gases, eight types of colloidal systems are possible. A gas mixed with another gas forms a
homogeneous mixture and hence is not a colloidal system. The examples of the various
types of colloids along with their typical names are listed in Table 5.4.

5.4.2 Classification Based on Nature of Interaction between Dispersed Phase and
Dispersion Medium

Depending upon the nature of interaction between the dispersed phase and the dispersion
medium, colloidal sols are divided into two categories, namely, lyophilic (solvent attracting) and
lyophobic (solvent repelling). If water is the dispersion medium, the terms used are
hydrophilic and hydrophobic.
(i) Lyophilic colloids: The word ‘lyophilic’ means liquid-loving. Colloidal sols directly formed
by mixing substances like gum, gelatine, starch, rubber, etc., with a suitable liquid (the
dispersion medium) are called lyophilic sols. An important characteristic of these sols
is that if the dispersion medium is separated from the dispersed phase (say by
evaporation), the sol can be reconstituted by simply remixing with the dispersion
medium. That is why these sols are also called reversible sols. Furthermore, these
sols are quite stable and cannot be easily coagulated as discussed later.
(ii) Lyophobic colloids: The word ‘lyophobic’ means liquid-hating. Substances like metals,
their sulphides, etc., when simply mixed with the dispersion medium do not form the
colloidal sol. Their colloidal sols can be prepared only by special methods (as discussed
later). Such sols are called lyophobic sols. These sols are readily precipitated (or
coagulated) on the addition of small amounts of electrolytes, by heating or by shaking and
hence, are not stable. Further, once precipitated, they do not give back the colloidal sol
by simple addition of the dispersion medium. Hence, these sols are also called
irreversible sols. Lyophobic sols need stabilising agents for their preservation.

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 51 of 74

m
m .co

m .co s em
se g la
(iii) Charge on colloidal particles: Colloidal particles aalways carry an electric charge.
The nature of this charge is the same on all the particles in a given colloidal solution
and may be either positive or negative. A list of some common sols with the nature of
charge on their particles is given below:

m
m .co
m .co s e m
s e l a
g l a ag
a
Electrophoresis: The existence of charge on colloidal particles is confirmed by electrophoresis
experiment. When electric potential is applied across two platinum electrodes dipping in a
colloidal solution, the colloidal particles move towards one or the other electrode. The movement
of colloidal particles under an applied electric potential is called electrophoresis. Positively
charged particles move towards the cathode while negatively charged particles move
towards the anode. This can be demonstrated by the following experimental set- up (Fig. 5.14).

m
m .co
s e
g la
a

m
m .co
.co s e m
s emWhen electrophoresis, i.e., movement of particles is prevented by gsome
l a suitable means, it is
g la observed that the dispersion medium begins to move in an electrica field. This phenomenon is
a termed electroosmosis

Coagulation or precipitation: The stability of the lyophobic sols is due to the presence of
charge on colloidal particles. If, somehow, the charge is removed, the particles will come nearer
to each other to form aggregates (or coagulate) and settle down under the force of gravity.
The process of settling of colloidal particles is called coagulation or
precipitation of the sol.
The coagulation of the lyophobic sols can be carried out in the following ways:

m
(i) By electrophoresis: The colloidal particles move towards oppositely charged electrodes, get
.
.co e m
discharged and precipitated.

em l as
las ag
ag For more Question Papers, Sample Papers, Notes & Syllabus visit Page 52 of 74

(ii) By mixing two oppositely charged sols: Oppositely charged sols when mixed in almost equal
proportions, neutralise their charges and get partially or completely precipitated. Mixing of
hydrated ferric oxide (+ve sol) and arsenious sulphide (–ve sol) bring them in the precipitated
forms. This type of coagulation is called mutual coagulation.
(iii) By boiling: When a sol is boiled, the adsorbed layer is disturbed due to increased collisions
with the molecules of dispersion medium. This reduces the charge on the particles and
ultimately leads to settling down in the form of a precipitate.
(iv) By persistent dialysis: On prolonged dialysis, traces of the electrolyte present in the sol
are removed almost completely and the colloids become unstable and ultimately
coagulate.
(v) By addition of electrolytes: When excess of an electrolyte is added, the colloidal particles
are precipitated. The reason is that colloids interact with ions carrying charge opposite to
that present on themselves. This causes neutralisation leading to their coagulation. The ion
responsible for neutralisation of charge on the particles is called the coagulating ion. A
negative ion causes the precipitation of positively charged sol and vice versa.

It has been observed that, generally, the greater the valence of the flocculating ion added,
the greater is its power to cause precipitation. This is known as Hardy-Schulze rule. In the
coagulation of a negative sol, the flocculating power is in the order: Al 3+>Ba2+>Na+
Similarly, in the coagulation of a positive sol, the flocculating power is in the order:
[Fe(CN)6]4– > PO 43– > SO 42– > Cl–
The minimum concentration of an electrolyte in millimoles per litre required to cause
precipitation of a sol in two hours is called coagulating value. The smaller the quantity needed,
the higher will be the coagulating power of an ion.

Coagulation of lyophilic sols

There are two factors which are responsible for the stability of lyophilic sols. These factors are
the charge and solvation of the colloidal particles. When these two factors are removed, a
lyophilic sol can be coagulated. This is done (i) by adding an electrolyte and (ii) by adding a
suitable solvent. When solvents such as alcohol and acetone are added to hydrophilic sols,
the dehydration of dispersed phase occurs. Under this condition, a small quantity of electrolyte
can bring about coagulation.
Protection of colloids

Lyophilic sols are more stable than lyophobic sols. This is due to the fact that lyophilic colloids
are extensively solvated, i.e., colloidal particles are covered by a sheath of the liquid in which
they are dispersed.
Lyophilic colloids have a unique property of protecting lyophobic colloids. When a lyophilic sol is
added to the lyophobic sol, the lyophilic particles form a layer around lyophobic particles and thus
protect the latter from electrolytes. Lyophilic colloids used for this purpose are called protective
colloids.

5.5 Emulsions
These are liquid-liquid colloidal systems, i.e., the dispersion of finely divided droplets in
another liquid. If a mixture of two immiscible or partially miscible liquids is shaken, a coarse
dispersion of one liquid in the other is obtained which is called emulsion. Generally, one of
the two liquids is water. There are two types of emulsions.
(i) Oil dispersed in water (O/W type) and
(ii) Water dispersed in oil (W/O type).

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 53 of 74

In the first system, water acts as dispersion medium. Examples of this type of emulsion are
milk and vanishing cream. In milk, liquid fat is dispersed in water. In the second system, oil
acts as dispersion medium.

Common examples of this type are butter and cream.
Emulsions of oil in water are unstable and sometimes they separate into two layers on
standing. For stabilisation of an emulsion, a third component called emulsifying agent is
usually added. The emulsifying agent forms an interfacial film between suspended particles
and the medium. The principal emulsifying agents for O/W emulsions are proteins, gums,
natural and synthetic soaps, etc., and for W/O, heavy metal salts of fatty acids, long chain
alcohols, lampblack, etc.

Emulsions can be diluted with any amount of the dispersion medium. On the other hand, the
dispersed liquid when mixed, forms a separate layer. The droplets in emulsions are often
negatively charged and can be precipitated by electrolytes. They also show Brownian
movement and Tyndall effect. Emulsions can be broken into constituent liquids by heating,
freezing, centrifuging, etc.

5.6 Colloids Around Us
Most of the substances, we come across in our daily life, are colloids. The meals we eat, the
clothes we wear, the wooden furniture we use, the houses we live in, the newspapers we
read, are largely composed of colloids.
Following are the interesting and noteworthy examples of colloids:
(i) Blue colour of the sky: Dust particles along with water suspended in air scatter blue light which
reaches our eyes and the sky looks blue to us.
(ii) Fog, mist and rain: When a large mass of air containing dust particles, is cooled below its
dewpoint, the moisture from the air condenses on the surfaces of these particles forming
fine droplets. These droplets being colloidal in nature continue to float in air in the form of
mist or fog. Clouds are aerosols having small droplets of water suspended in air. On
account of condensation in the upper atmosphere, the colloidal droplets of water grow
bigger and bigger in size, till they come down in the form of rain. Sometimes, the rainfall
occurs when two oppositely charged clouds meet.
It is possible to cause artificial rain by throwing electrified sand or spraying a sol carrying
charge opposite to the one on clouds from an aeroplane.
(iii) Food articles: Milk, butter, halwa, ice creams, fruit juices, etc., are all colloids in one form
or the other.
(iv) Blood: It is a colloidal solution of an albuminoid substance. The styptic action of alum and
ferric chloride solution is due to coagulation of blood forming a clot which stops further
bleeding.
(v) Soils: Fertile soils are colloidal in nature in which humus acts as a protective colloid. On
account of colloidal nature, soils adsorb moisture and nourishing materials.
(vi) Formation of delta: River water is a colloidal solution of clay. Sea water contains a number of
electrolytes. When river water meets the sea water, the electrolytes present in sea water
coagulate the colloidal solution of clay resulting in its deposition with the formation of delta.

****************************************************************************************************************

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 54 of 74

m
m .co

m .co s em
se g la
a
GENERAL PRINCIPLES AND PROCESSES OF ISOLATION OF
ELEMENTS
The history of civilisation is linked to the use of metals in antiquity in many ways. Different
periods of early human civilisations have been named after metals. The skill of extraction of
metals gave many metals and brought about several changes in the human society. It gave
weapons, tools, ornaments, utensils, etc., and enriched the cultural life. The ‘Seven metals
of antiquity’, as they are sometimes called, are gold, copper, silver, lead, tin, iron and
mercury. Although modern metallurgy had exponential growth after Industrial Revolution, it
m
.co
is interesting to note that many modern concepts in metallurgy have their roots in ancient

m
practices that pre-dated the Industrial Revolution. For over 7000 years, India has had a rich

.co m
tradition of metallurgical skills.
The two important sources for the history of Indian metallurgy are archaeological excavations
m s e
e
and literary evidences. The first evidence of metal in Indian subcontinent comes from
s l a
a ag
Mehrgarh in Baluchistan, where a small copper bead, dated to about 6000 BCE was found.

g l
It is however thought to be native copper, which has not been extracted from the ore.

a
Spectrometric studies on copper ore samples obtained from the ancient mine pits at Khetri in
Rajasthan and on metal samples cut from representative Harappan artefacts recovered from
Mitathal in Haryana and eight other sites distributed in Rajasthan, Gujarat, Madhya Pradesh
and Maharashtra prove that copper metallurgy in India dates back to the Chalcolithic
cultures in the subcontinent. Indian chalcolithic copper objects were in all probability made
indigenously. The ore for extraction of metal for making the objects was obtained from
chalcopyrite ore deposits in Aravalli Hills. Collection of archaeological texts from copper-
plates and rock-inscriptions have been compiled and published by the Archaeological
Survey of India during the past century. Royal records were engraved on copper plates
(tamra-patra). Earliest known copper- plate has a Mauryan record that mentions famine relief

m
efforts. It has one of the very few pre-Ashoka Brahmi inscriptions in India.

.co
Harappans also used gold and silver, as well as their joint alloy electrum. Variety of

m
ornaments such as pendants, bangles, beads and rings have been found in ceramic or

e
bronze pots. Early gold and silver ornaments have been found from Indus Valley sites such
s
la
as Mohenjodaro (3000 BCE). These are on display in the National Museum, New Delhi. India

g
has the distinction of having the deepest ancient gold mines in the world, in the Maski region
a
of Karnataka. Carbon dating places them in mid- 1st millennium BCE.
Hymns of Rigveda give earliest indirect references to the alluvial placer gold deposits in
India. The river Sindhu was an important source of gold in ancient times. It is interesting that
the availability of alluvial placer gold in the river Sindhu has been reported in modern times
also. It has been reported that there are great mines of gold in the region of Mansarovar and
in Thokjalyug even now. The Pali text, Anguttara Nikaya narrates the process of the recovery of
gold dust or particles from alluvial placer gold deposits. Although evidence of gold refining is
available in Vedic texts, it is Kautilya’s Arthashastra, authored probably in 3rd or 4th century
BCE, during Mauryan era, which has much data on prevailing chemical practices in a long
section on mines and minerals including metal ores of gold, silver, copper, lead, tin and iron.
Kautilya describes a variety of gold called rasviddha, which is naturally occurring gold
m
m .co
solution. Kalidas also mentioned about such solutions. It is astonishing how people

.co m
recognised such solutions.

m s e
The native gold has different colours depending upon the nature and amount of impurity present

s e l a
in it. It is likely that the different colours of native gold were a major driving force for the
g
la
development of gold refining.

g a
Recent excavations in central parts of Ganges Valley and Vindhya hills have shown that iron
a was produced there possibly as early as in 1800 BCE. In the recent excavations conducted
by the Uttar Pradesh State Archaeological Department, iron furnaces, artefacts, tuyers and
layers of slag have been found. Radiocarbon dating places them between BCE 1800 and 1000.
The results of excavation indicate that the knowledge of iron smelting and manufacturing of
iron artefacts was well known in Eastern Vindhyas and it was in use in the Central Ganga
Plains, at least from the early 2nd millennium BCE. The quantity and types of iron artefacts
and the level of technical advancements indicate that working of iron would have been
introduced much earlier.

The evidence indicates early use of iron in other areas of the country, which proves that India
m .
.co
was indeed an independent centre for the development of the working of iron.
Iron smelting and the use of iron was especially established in South Indian megalithic
e m
em l as
las ag
ag For more Question Papers, Sample Papers, Notes & Syllabus visit Page 55 of 74

cultures. The forging of wrought iron seems to have been at peak in India in the Ist
millennium CE. Greek accounts report the manufacture of steel in India by crucible process.
In this process, iron, charcoal and glass were mixed together in a crucible and heated until
the iron melted and absorbed the carbon. India was a major innovator in the production of
advanced quality steel. Indian steel was called ‘the Wonder Material of the Orient’. A Roman
historian, Quintus Curtius, records that one of the gifts Porus of Taxila (326 BCE) gave to
Alexander the Great was some two-and-a-half tons of Wootz steel. Wootz steel is primarily
iron containing a high proportion of carbon (1.0 – 1.9%). Wootz is the English version of the
word ‘ukku’ which is used for steel in Karnataka and Andhra Pradesh. Literary accounts
suggest that Indian Wootz steel from southern part of the Indian subcontinent was exported
to Europe, China and Arab world. It became prominent in the Middle East where it was
named as Damasus Steel. Michael Faraday tried to duplicate this steel by alloying iron with
a variety of metals, including noble metals, but failed.
When iron ore is reduced in solid state by using charcoal, porous iron blocks are formed.
Therefore, reduced iron blocks are also called sponge iron blocks. Any useful product can be
obtained from this material only after removing the porosity by hot forging. The iron so obtained
is termed as wrought iron. An exciting example of wrought iron produced in ancient India is the
world famous Iron Pillar. It was erected in its present position in Delhi in 5th century CE. The
Sanskrit inscription engraved on it suggests that it was brought here from elsewhere during the
Gupta Period. The average composition (weight%) of the components present in the wrought
iron of the pillar, besides iron, are 0.15% C, 0.05% Si, 0.05% Mn, 0.25% P, 0.005% Ni,
0.03%
Cu and 0.02% N. The most significant aspect of the pillar is that there is no sign of corrosion
inspite of the fact that it has been exposed to the atmosphere for about 1,600 years.
Radiocarbon dating of charcoal from iron slag revealed evidence of continuous smelting
in Khasi Hills of Meghalaya. The slag layer, which is dated to 353 BCE – CE 128, indicates
that Khasi Hill region is the earliest iron smelting site studied in the entire region of North
East India. The remnants of former iron-ore excavation and iron manufacturing are visible
even now in the landscape of Khasi Hills. British naturalists who visited Meghalaya in early
19th century described the iron industry that had developed in the upper part of the Khasi
Hills.
There is archaeological evidence of zinc production in Rajasthan mines at Zawar from the
6th or 5th BCE. India was the first country to master zinc distillation. Due to low boiling
point, zinc tends to vapourise while its ore is smelted. Pure zinc could be produced after a
sophisticated ‘downward’ distillation technique in which the vapour was condensed in a
lower container. This technique was also applied to mercury. Indian metallurgists were
masters in this technique. This has been described in Sanskrit texts of 14th century.
Indians had knowledge about mercury. They used it for medicinal purpose. Development of
mining and metallurgy declined during the British colonial era. By the 19th century, once flourished
mines of Rajasthan were mostly abandoned and became almost extinct. In 1947 when India got
independence, European literature on science had already found its way slowly into the country.
Thus, in post independence era, the Government of India initiated the process of nation building
through the establishment of various institutes of science and technology. In the following sections,
we will learn about the modern methods of extraction of elements.

6.1 Occurrence of Metals
A few elements like carbon, sulphur, gold and noble gases, occur in free state while others are
found in combined forms in the earth’s crust. Elements vary in abundance. Among metals,
aluminium is the most abundant. In fact, it is the third most abundant element in earth’s crust
(8.3% approx. by weight). It is a major component of many igneous minerals including mica
and clays. Many gemstones are impure forms of Al2O3. For example, gems ‘ruby’ and
‘sapphire’ have Cr and Co respectively as impurity. Iron is the second most abundant metal in
the earth’s crust. It forms a variety of compounds and their various uses make it a very important
element. It is one of the essential elements in biological systems as well.
For obtaining a particular metal, first we look for minerals which are naturally occurring
chemical substances in the earth’s crust and are obtained through mining. Out of many
minerals in which a metal may be found, only a few are viable to be used as source of that
metal. Such minerals are known as ores.
The principal ores of aluminium, iron, copper and zinc are given in Table 6.1.

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 56 of 74

A particular element may occur in a variety of compounds. The process of isolation of
element from its compound should be such that it is chemically feasible and commercially
viable. For the purpose of extraction, bauxite is chosen for aluminium. For iron, usually the
oxide ores which are abundant and do not produce polluting gases (like SO 2 that is produced
in case of iron pyrites) are taken. For copper and zinc, any of the ores listed in Table 6.1
may be used depending upon the availability and other relevant factors.
The entire scientific and technological process used for isolation of the metal from its ore is
known as metallurgy. The extraction and isolation of an element from its combined form
involves various principles of chemistry. Still, some general principles are common to all the
extraction processes of metals.
An ore rarely contains only a desired substance. It is usually contaminated with earthly or
undesired materials known as gangue. The extraction and isolation of metals from ores
involves the following major steps:
• Concentration of the ore,
• Isolation of the metal from its concentrated ore, and
• Purification of the metal.
In the following Sections, we shall first describe the various steps for effective
concentration of ores. After that principles of some of the common metallurgical processes
will be discussed. Those principles will include the thermodynamic and electrochemical
aspects involved in the effective reduction of the concentrated ore to the metal.

6.2 Concentration of Ores
Removal of the unwanted materials (e.g., sand, clays, etc.) from the ore is known as
concentration, dressing or benefaction. Before proceeding for concentration, ores are
graded and crushed to reasonable size. Concentration of ores involves several steps and
selection of these steps
depends upon the differences in physical properties of the compound of the metal present and
that of the gangue. The type of the metal, the available facilities and the environmental factors
are also taken into consideration. Some of the important procedures for concentration of ore are
described below.

6.2.1 Hydraulic Washing
This is based on the difference between specific gravities of the ore and the gangue particles.
It is therefore a type of gravity separation. In one such process, an upward stream of running
water is used to wash the powdered ore. The lighter gangue particles are washed away and
the heavier ore particles are left behind.

6.2.2 Magnetic Separation

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 57 of 74

m
m .co

m .co s em
se g la
a components. If either the ore or
This is based on differences in magnetic properties of the ore
the gangue is attracted towards magnetic field, then the separation is carried out by this
method. For example iron ores are attracted towards magnet, hence, non–magnetic impurities
can be separted from them using magnetic separation. The powdered ore is dropped over a
conveyer belt which moves over a magnetic roller (Fig.6.1) Magnetic substance remains
attracted towards the belt and falls close to it.

m
m .co
.co s e m
s em l a
g la Fig. 6.1: Magnetic separation (schematic) ag
a Floatation Method
6.2.3 Froth
This method is used for removing gangue from sulphide ores. In this process, a suspension
of the powdered ore is made with water. Collectors and froth stabilisers are added to it.
Collectors (e.g., pine oils, fatty acids, xanthates, etc.) enhance non- wettability of the mineral
particles and froth stabilisers (e.g., cresols, aniline) stabilise the froth. The mineral particles
become wet by oils while the gangue particles by water. A rotating paddle agitates the mixture
and draws a i r i n i t . As a result, froth is formed which carries the mineral particles. The
froth is light and skimmed off. It is then dried for recovery of the ore particles. Sometimes,
it is possible to separate two sulphide ores by adjusting p r o p o r t i o n o f o i l t o w a t e r
m
.co
o r b y u s i n g ‘ depressants’.
For example, in the case of an ore containing ZnS and PbS, the depressant used is
NaCN. It selectively prevents ZnS from coming
e m
las
to the froth but allows PbS to come with the froth.

ag

m
m .co
.co s e m
s em Fig. 6.2: Froth floatation process
g l a
g la The Innovative Washerwoman a
a One can do wonders if he or she has a scientific temperament and is attentive to
observations. A washerwoman had an innovative mind too. While washing a miner’s
overalls, she noticed that sand and similar dirt fell to the bottom of the washtub. What was
peculiar, the copper bearing compounds that had come to the clothes from the mines, were
caught in the soapsuds and so they came to the top. One of her clients, Mrs. Carrie Everson
was a chemist. The washerwoman told her experience to Mrs. Everson. The latter thought
that the idea could be used for separating copper compounds from rocky and earth materials
on large scale. This way an invention came up. At that time only those ores were used for
m .
.co
extraction of copper, which contained large amounts of the metal. Invention of the Froth

e m
em l as
las ag
ag For more Question Papers, Sample Papers, Notes & Syllabus visit Page 58 of 74

Floatation Method made copper mining profitable even from the low- grade ores. World
production of copper soared and the metal became cheaper.

6.2.4 Leaching
Leaching is often used if the ore is soluble in some suitable solvent. Following examples
illustrate the procedure:

(a) Leaching of alumina from bauxite
Bauxite is the principal ore of aluminium. It usually contains SiO2, iron oxides and titanium
oxide (TiO 2) as impurities. Concentration is carried out by heating the powdered ore with a
concentrated solution of NaOH at 473 – 523 K and 35 – 36 bar pressure. This process is
called digestion. This way, Al2O3 is extracted out as sodium aluminate.

The impurity, SiO 2 too dissolves forming sodium silicate. Other impurities are left behind.
Al2O3(s) + 2NaOH(aq) + 3H2O(l)  Na[Al(OH)4](aq) (6.1)

The sodium aluminate present in solution is neutralised by passing CO 2 gas and hydrated
Al2O3 is precipitated. At this stage, small amount of freshly prepared sample of hydrated
Al2O3 is added to the solution. This is called seeding. It induces the precipitation.
2Na[Al(OH)4](aq) + 2CO2(g)  Al2O3.xH2O(s) + 2NaHCO3 (aq) (6.2)
Sodium silicate remains in the solution and hydrated alumina is filtered, dried and heated to
give back pure Al2O3.

(b) Other example

In the metallurgy of silver and gold, the respective metal is leached with a dilute solution of
NaCN or KCN in the presence of air, which supplies O 2. The metal is obtained later by
replacement reaction.
4M(s) + 8CN- (aq) + 2H2O(aq) + O2(g)  4[M(CN)2]-(aq) + 4OH-(aq) (M=Ag or Au) (6.4)

2[M(CN)2]-(aq) + Zn(s) {Zn(CN)4]2- (aq) + 2M(s) (6.5)

6.3 Extractionof Crude Metal from Concentrated Ore
To extract metal from concentrated ore, it must be converted to a form which is suitable for
reduction to metal. Usually sulphide ores are converted to oxide before reduction because
oxides are easier to reduce. Thus isolation of metals from concentrated ore involves two major
steps viz.,
(a) conversion to oxide, and
(b) reduction of the oxide to metal
(a) Conversion t o o x i d e

(i) Calcination: Calcinaton involves heating. It removes the volatile matter which escapes
leaving behind the metal oxide:

For more Question Papers, Sample Papers, Notes & Syllabus visit Page 59 of 74