Chemistry Class 11 Part 2
Chemistry Class 11 Part 2
Chemistry Class 11 Part 2
CONTENTS
FOREWORD iii
Unit 8 Redox Reactions 255
8.1 Classical Idea of Redox Reactions-Oxidation and Reduction Reactions 255
8.2 Redox Reactions in Terms of Electron Transfer Reactions 257
8.3 Oxidation Number 259
8.4 Redox Reactions and Electrode Processes 269
UNIT 9
HYDROGEN
possess metallic characteristics under normal solar atmosphere. The giant planets Jupiter
conditions. In fact, in terms of ionization and Saturn consist mostly of hydrogen.
enthalpy, hydrogen resembles more However, due to its light nature, it is much less
–1
with halogens, Δi H of Li is 520 kJ mol , F is abundant (0.15% by mass) in the earth’s
–1 –1
1680 kJ mol and that of H is 1312 kJ mol . atmosphere. Of course, in the combined form
Like halogens, it forms a diatomic molecule, it constitutes 15.4% of the earth's crust and
combines with elements to form hydrides and the oceans. In the combined form besides in
a large number of covalent compounds. water, it occurs in plant and animal tissues,
However, in terms of reactivity, it is very low as carbohydrates, proteins, hydrides including
compared to halogens. hydrocarbons and many other compounds.
Inspite of the fact that hydrogen, to a 9.2.2 Isotopes of Hydrogen
certain extent resembles both with alkali Hydrogen has three isotopes: protium, 1 H,
1
high H–H bond enthalpy. Thus, the atomic (i) Hydrogenation of vegetable oils using
hydrogen is produced at a high temperature nickel as catalyst gives edible fats
in an electric arc or under ultraviolet (margarine and vanaspati ghee)
radiations. Since its orbital is incomplete with (ii) Hydroformylation of olefins yields
1
1s electronic configuration, it does combine aldehydes which further undergo
with almost all the elements. It accomplishes reduction to give alcohols.
reactions by (i) loss of the only electron to
+
give H , (ii) gain of an electron to form H , and
–
H2 + CO + RCH = CH2 → RCH 2CH2CHO
(iii) sharing electrons to form a single covalent bond.
H 2 + RCH 2CH 2 CHO → RCH 2 CH 2CH2 OH
The chemistry of dihydrogen can be
illustrated by the following reactions: Problem 9.1
Reaction with halogens: It reacts with Comment on the reactions of dihydrogen
halogens, X2 to give hydrogen halides, HX, with (i) chlorine, (ii) sodium, and (iii)
H2 ( g ) + X 2 ( g ) → 2HX ( g ) (X = F,Cl, Br,I) copper(II) oxide
Solution
While the reaction with fluorine occurs even in
the dark, with iodine it requires a catalyst. (i) Dihydrogen reduces chlorine into
–
chloride (Cl ) ion and itself gets oxidised
Reaction with dioxygen: It reacts with +
to H ion by chlorine to form hydrogen
dioxygen to form water. The reaction is highly
chloride. An electron pair is shared
exothermic.
between H and Cl leading to the formation
2H2(g) + O2 (g) 2H2O(l); of a covalent molecule.
ΔH = –285.9 kJ mol
V –1
(ii) Dihydrogen is reduced by sodium to
Reaction with dinitrogen: With dinitrogen form NaH. An electron is transferred from
it forms ammonia. Na to H leading to the formation of an ionic
+ –
compound, Na H .
3H2 ( g ) + N 2 ( g ) ⎯⎯⎯⎯⎯⎯
673K,200atm
Fe
→ 2NH3 ( g ) ; (iii) Dihydrogen reduces copper(II) oxide
ΔH V = −92.6 kJ mol −1 to copper in zero oxidation state and itself
gets oxidised to H2O, which is a covalent
This is the method for the manufacture of molecule.
ammonia by the Haber process.
Reactions with metals: With many metals it 9.4.3 Uses of Dihydrogen
combines at high a temperature to yield the • The largest single use of dihydrogen is in
corresponding hydrides (section 9.5) the synthesis of ammonia which is used in
H2(g) +2M(g) → 2MH(s); the manufacture of nitric acid and
where M is an alkali metal nitrogenous fertilizers.
Reactions with metal ions and metal • Dihydrogen is used in the manufacture of
oxides: It reduces some metal ions in aqueous vanaspati fat by the hydrogenation of
solution and oxides of metals (less active than polyunsaturated vegetable oils like
iron) into corresponding metals. soyabean, cotton seeds etc.
• It is used in the manufacture of bulk
H2 ( g ) + Pd2 + ( aq ) → Pd ( s ) + 2H + ( aq )
organic chemicals, particularly methanol.
yH2 ( g ) + M x O y ( s ) → xM ( s ) + yH2O ( l )
CO ( g ) + 2H2 ( g ) ⎯⎯⎯⎯
cobalt
catalyst
→ CH3OH ( l )
Reactions with organic compounds: It
• It is widely used for the manufacture of
reacts with many organic compounds in the
presence of catalysts to give useful metal hydrides (section 9.5)
hydrogenated products of commercial • It is used for the preparation of hydrogen
importance. For example : chloride, a highly useful chemical.
280 CHEMISTRY
Solution Solution
On the basis of molecular masses of NH3, Although phosphorus exhibits +3 and +5
H2O and HF, their boiling points are oxidation states, it cannot form PH5.
expected to be lower than those of the Besides some other considerations, high
subsequent group member hydrides. ΔaH value of dihydrogen and ΔegH value
However, due to higher electronegativity of hydrogen do not favour to exhibit the
of N, O and F, the magnitude of hydrogen highest oxidation state of P, and
bonding in their hydrides will be quite consequently the formation of PH5.
appreciable. Hence, the boiling points
NH3, H2O and HF will be higher than the
hydrides of their subsequent group 9.6 WATER
members. A major part of all living organisms is made
up of water. Human body has about 65% and
9.5.3 Metallic or Non-stoichiometric some plants have as much as 95% water. It is
(or Interstitial ) Hydrides a crucial compound for the survival of all life
These are formed by many d-block and f-block forms. It is a solvent of great importance. The
elements. However, the metals of group 7, 8 distribution of water over the earth’s surface
and 9 do not form hydride. Even from group is not uniform. The estimated world water
6, only chromium forms CrH. These hydrides supply is given in Table 9.2
conduct heat and electricity though not as Table 9.2 Estimated World Water Supply
efficiently as their parent metals do. Unlike
saline hydrides, they are almost always non- Source % of Total
stoichiometric, being deficient in hydrogen. For
Oceans 97.33
example, LaH2.87, YbH2.55, TiH1.5–1.8, ZrH1.3–1.75,
Saline lakes and inland seas 0.008
VH 0.56 , NiH 0.6–0.7 , PdH 0.6–0.8 etc. In such
hydrides, the law of constant composition does Polar ice and glaciers 2.04
not hold good. Ground water 0.61
Earlier it was thought that in these Lakes 0.009
hydrides, hydrogen occupies interstices in the Soil moisture 0.005
metal lattice producing distortion without any Atmospheric water vapour 0.001
change in its type. Consequently, they were Rivers 0.0001
termed as interstitial hydrides. However, recent
studies have shown that except for hydrides 9.6.1 Physical Properties of Water
of Ni, Pd, Ce and Ac, other hydrides of this class It is a colourless and tasteless liquid. Its
have lattice different from that of the parent physical properties are given in Table 9.3 along
metal. The property of absorption of hydrogen with the physical properties of heavy water.
on transition metals is widely used in catalytic The unusual properties of water in the
reduction / hydrogenation reactions for the
condensed phase (liquid and solid states) are
preparation of large number of compounds.
due to the presence of extensive hydrogen
Some of the metals (e.g., Pd, Pt) can
bonding between water molecules. This leads
accommodate a very large volume of hydrogen
to high freezing point, high boiling point, high
and, therefore, can be used as its storage
heat of vaporisation and high heat of fusion in
media. This property has high potential for
comparison to H2S and H2Se. In comparison
hydrogen storage and as a source of energy.
to other liquids, water has a higher specific
heat, thermal conductivity, surface tension,
Problem 9.3
dipole moment and dielectric constant, etc.
Can phosphorus with outer electronic These properties allow water to play a key role
2 3
configuration 3s 3p form PH5 ?
in the biosphere.
282 CHEMISTRY
The high heat of vaporisation and heat polar molecule, (Fig 9.1(b)). Its orbital overlap
capacity are responsible for moderation of the picture is shown in Fig. 9.1(c). In the liquid
climate and body temperature of living beings. phase water molecules are associated together
It is an excellent solvent for transportation of by hydrogen bonds.
ions and molecules required for plant and The crystalline form of water is ice. At
animal metabolism. Due to hydrogen bonding atmospheric pressure ice crystallises in the
with polar molecules, even covalent hexagonal form, but at very low temperatures
compounds like alcohol and carbohydrates it condenses to cubic form. Density of ice is
dissolve in water. less than that of water. Therefore, an ice cube
9.6.2 Structure of Water floats on water. In winter season ice formed
on the surface of a lake provides thermal
In the gas phase water is a bent molecule with
insulation which ensures the survival of the
a bond angle of 104.5°, and O–H bond length
aquatic life. This fact is of great ecological
of 95.7 pm as shown in Fig 9.1(a). It is a highly
significance.
9.6.3 Structure of Ice
Ice has a highly ordered three dimensional
hydrogen bonded structure as shown in
Fig. 9.2. Examination of ice crystals with
RNH .OH+
3
–
( s ) + X ( aq ) U RNH
− +
3 (s)
.X −
( oxidised product )
+ OH − ( aq ) In this case 1% H 2O2 is formed. It is
– – – 2– extracted with water and concentrated to ~30%
OH exchanges for anions like Cl , HCO3, SO4 (by mass) by distillation under reduced
–
etc. present in water. OH ions, thus, liberated pressure. It can be further concentrated to
+
neutralise the H ions set free in the cation ~85% by careful distillation under low
exchange. pressure. The remaining water can be frozen
H+ ( aq ) + OH − ( aq ) → H2 O ( l ) out to obtain pure H2O2.
The exhausted cation and anion exchange 9.7.2 Physical Properties
resin beds are regenerated by treatment with In the pure state H2O2 is an almost colourless
dilute acid and alkali solutions respectively. (very pale blue) liquid. Its important physical
properties are given in Table 9.4.
9.7 HYDROGEN PEROXIDE (H2O2)
H 2 O 2 is miscible with water in all
Hydrogen peroxide is an important chemical
proportions and forms a hydrate H2O2.H2O
used in pollution control treatment of domestic
(mp 221K). A 30% solution of H2O2 is marketed
and industrial effluents.
as ‘100 volume’ hydrogen peroxide. It means that
9.7.1 Preparation one millilitre of 30% H2O2 solution will give 100 V
It can be prepared by the following methods. of oxygen at STP. Commercially, it is marketed
(i) Acidifying barium peroxide and removing as 10 V, which means it contains 3% H2O2.
excess water by evaporation under reduced Problem 9.4
pressure gives hydrogen peroxide.
Calculate the strength of 10 volume
BaO2 .8H2 O ( s ) + H2 SO4 ( aq ) → BaSO4 ( s ) + solution of hydrogen peroxide.
H2 O2 ( aq ) + 8H2O ( l ) Solution
10 volume solution of H2O2 means that
(ii) Peroxodisulphate, obtained by electrolytic 1L of this H2O2 will give 10 L of oxygen at STP
oxidation of acidified sulphate solutions at
high current density, on hydrolysis yields 2H2O2 ( l ) → O2 ( g ) + H2O ( l )
hydrogen peroxide. 2×34 g 22.4 L at STP
68 g
2HSO4− ( aq ) ⎯⎯⎯⎯⎯
Electrolysis
→ HO3 SOOSO3 H ( aq )
22.4 L of O2 at STP is produced from
⎯⎯⎯⎯⎯
Hydrolysis
→ 2HSO4− ( aq ) + 2H+ ( aq ) + H2O2 ( aq )
H2O2 = 68 g
This method is now used for the laboratory 10 L of O2 at STP is produced from
preparation of D2O2. 68 × 10
H2 O2 = g = 30.36 g
K2S2O8 ( s) + 2D2O ( l ) → 2KDSO4 ( aq ) + D2O2 ( l ) 22.4
Therefore, strength of H2O2 in 10 volume
(iii) Industrially it is prepared by the auto-
H2O2 = 30.36 g/L
oxidation of 2-alklylanthraquinols.
Table 9.4 Physical Properties of Hydrogen Peroxide
–3
Melting point/K 272.4 Density (liquid at 298 K)/g cm 1.44
Boiling point(exrapolated)/K 423 Viscosity (290K)/centipoise 1.25
2 2
Vapour pressure(298K)/mmHg 1.9 Dielectric constant (298K)/C /N m 70.7
–3 –1 –1 –8
Density (solid at 268.5K)/g cm 1.64 Electrical conductivity (298K)/Ω cm 5.1×10
286 CHEMISTRY
amounts in mole, mass and volume, are shown limitations have prompted researchers to
in Table 9.3. search for alternative techniques to use
From this table it is clear that on a mass dihydrogen in an efficient way.
for mass basis dihydrogen can release more In this view Hydrogen Economy is an
energy than petrol (about three times). alternative. The basic principle of hydrogen
Moreover, pollutants in combustion of economy is the transportation and storage of
dihydrogen will be less than petrol. The only energy in the form of liquid or gaseous
pollutants will be the oxides of dinitrogen (due dihydrogen. Advantage of hydrogen economy
to the presence of dinitrogen as impurity with is that energy is transmitted in the form of
dihydrogen). This, of course, can be minimised dihydrogen and not as electric power. It is for
by injecting a small amount of water into the the first time in the history of India that a pilot
cylinder to lower the temperature so that the project using dihydrogen as fuel was launched
reaction between dinitrogen and dioxygen may in October 2005 for running automobiles.
not take place. However, the mass of the Initially 5% dihydrogen has been mixed in
containers in which dihydrogen will be kept CNG for use in four-wheeler vehicles. The
must be taken into consideration. A cylinder percentage of dihydrogen would be gradually
of compressed dihydrogen weighs about 30 increased to reach the optimum level.
times as much as a tank of petrol containing
the same amount of energy. Also, dihydrogen Nowadays, it is also used in fuel cells for
gas is converted into liquid state by cooling to generation of electric power. It is expected that
20K. This would require expensive insulated economically viable and safe sources of
tanks. Tanks of metal alloy like NaNi5, Ti–TiH2, dihydrogen will be identified in the years to
Mg–MgH 2 etc. are in use for storage of come, for its usage as a common source of
dihydrogen in small quantities. These energy.
Table 9.3 The Energy Released by Combustion of Various Fuels in Moles, Mass and Volume
SUMMARY
Hydrogen is the lightest atom with only one electron. Loss of this electron results in an
elementary particle, the proton. Thus, it is unique in character. It has three isotopes,
1 2 3
namely : protium (1H), deuterium (D or 1H) and tritium (T or 1H). Amongst these three,
only tritium is radioactive. Inspite of its resemblance both with alkali metals and halogens,
it occupies a separate position in the periodic table because of its unique properties.
Hydrogen is the most abundant element in the universe. In the free state it is almost
not found in the earth’s atmosphere. However, in the combined state, it is the third most
abundant element on the earth’s surface.
Dihydrogen on the industrial scale is prepared by the water-gas shift reaction from
petrochemicals. It is obtained as a byproduct by the electrolysis of brine.
288 CHEMISTRY
–1
The H–H bond dissociation enthalpy of dihydrogen (435.88 kJ mol ) is the highest
for a single bond between two atoms of any elements. This property is made use of in the
atomic hydrogen torch which generates a temperature of ~4000K and is ideal for welding
of high melting metals.
Though dihydrogen is rather inactive at room temperature because of very high
negative dissociation enthalpy, it combines with almost all the elements under appropriate
conditions to form hydrides. All the type of hydrides can be classified into three categories:
ionic or saline hydrides, covalent or molecular hydrides and metallic or non-stoichiometric
hydrides. Alkali metal hydrides are good reagents for preparing other hydride compounds.
Molecular hydrides (e.g., B2H6, CH4, NH3, H2O) are of great importance in day-to-day life.
Metallic hydrides are useful for ultrapurification of dihydrogen and as dihydrogen storage
media.
Among the other chemical reactions of dihydrogen, reducing reactions leading to
the formation hydrogen halides, water, ammonia, methanol, vanaspati ghee, etc. are of
great importance. In metallurgical process, it is used to reduce metal oxides. In space
programmes, it is used as a rocket fuel. In fact, it has promising potential for use as a
non-polluting fuel of the near future (Hydrogen Economy).
Water is the most common and abundantly available substance. It is of a great
chemical and biological significance. The ease with which water is transformed from
liquid to solid and to gaseous state allows it to play a vital role in the biosphere. The
water molecule is highly polar in nature due to its bent structure. This property leads to
hydrogen bonding which is the maximum in ice and least in water vapour. The polar
nature of water makes it: (a) a very good solvent for ionic and partially ionic compounds;
(b) to act as an amphoteric (acid as well as base) substance; and (c) to form hydrates of
different types. Its property to dissolve many salts, particularly in large quantity, makes
it hard and hazardous for industrial use. Both temporary and permanent hardness can
be removed by the use of zeolites, and synthetic ion-exchangers.
Heavy water, D2O is another important compound which is manufactured by the
electrolytic enrichment of normal water. It is essentially used as a moderator in nuclear
reactors.
Hydrogen peroxide, H2O2 has an interesting non-polar structure and is widely used
as an industrial bleach and in pharmaceutical and pollution control treatment of
industrial and domestic effluents.
EXERCISES
9.1 Justify the position of hydrogen in the periodic table on the basis of its electronic
configuration.
9.2 Write the names of isotopes of hydrogen. What is the mass ratio of these isotopes?
9.3 Why does hydrogen occur in a diatomic form rather than in a monoatomic form
under normal conditions?
9.4 How can the production of dihydrogen, obtained from ‘coal gasification’, be
increased?
9.5 Describe the bulk preparation of dihydrogen by electrolytic method. What is the
role of an electrolyte in this process ?
9.6 Complete the following reactions:
Δ
(i) H 2 ( g ) + M m O o ( s ) ⎯⎯⎯→
(ii) CO ( g ) + H 2 ( g ) ⎯ ⎯ Δ
⎯ ⎯→
catalyst
HYDROGEN 289
Δ
(iii) C3H8 ( g ) + 3H 2O ( g ) ⎯⎯⎯⎯
catalyst
→
(vi) Ca 3 N 2 ( s ) + H 2O ( l ) →
Classify the above into (a) hydrolysis, (b) redox and (c) hydration reactions.
9.21 Describe the structure of the common form of ice.
9.22 What causes the temporary and permanent hardness of water ?
9.23 Discuss the principle and method of softening of hard water by synthetic ion-
exchange resins.
9.24 Write chemical reactions to show the amphoteric nature of water.
9.25 Write chemical reactions to justify that hydrogen peroxide can function as an
oxidising as well as reducing agent.
290 CHEMISTRY
UNIT 10
Lithium and beryllium, the first elements increase in atomic number, the atom becomes
+
of Group 1 and Group 2 respectively exhibit larger. The monovalent ions (M ) are smaller
some properties which are different from those than the parent atom. The atomic and ionic
of the other members of the respective group. radii of alkali metals increase on moving down
In these anomalous properties they resemble the group i.e., they increase in size while going
the second element of the following group. from Li to Cs.
Thus, lithium shows similarities to magnesium 10.1.3 Ionization Enthalpy
and beryllium to aluminium in many of their
The ionization enthalpies of the alkali metals
properties. This type of diagonal similarity is
are considerably low and decrease down the
commonly referred to as diagonal relationship
group from Li to Cs. This is because the effect
in the periodic table. The diagonal relationship
of increasing size outweighs the increasing
is due to the similarity in ionic sizes and /or
nuclear charge, and the outermost electron is
charge/radius ratio of the elements.
very well screened from the nuclear charge.
Monovalent sodium and potassium ions and
divalent magnesium and calcium ions are 10.1.4 Hydration Enthalpy
found in large proportions in biological fluids. The hydration enthalpies of alkali metal ions
These ions perform important biological decrease with increase in ionic sizes.
functions such as maintenance of ion balance + + + + +
Li > Na > K > Rb > Cs
and nerve impulse conduction. +
Li has maximum degree of hydration and
10.1 GROUP 1 ELEMENTS: ALKALI for this reason lithium salts are mostly
METALS hydrated, e.g., LiCl· 2H2O
The alkali metals show regular trends in their 10.1.5 Physical Properties
physical and chemical properties with the
All the alkali metals are silvery white, soft and
increasing atomic number. The atomic,
light metals. Because of the large size, these
physical and chemical properties of alkali
elements have low density which increases
metals are discussed below.
down the group from Li to Cs. However,
10.1.1 Electronic Configuration potassium is lighter than sodium. The melting
All the alkali metals have one valence electron, and boiling points of the alkali metals are low
1
ns (Table 10.1) outside the noble gas core. indicating weak metallic bonding due to the
The loosely held s-electron in the outermost presence of only a single valence electron in
valence shell of these elements makes them the them. The alkali metals and their salts impart
most electropositive metals. They readily lose characteristic colour to an oxidizing flame. This
+
electron to give monovalent M ions. Hence they is because the heat from the flame excites the
are never found in free state in nature. outermost orbital electron to a higher energy
level. When the excited electron comes back to
Element Symbol Electronic configuration the ground state, there is emission of radiation
in the visible region as given below:
Lithium Li 1s22s1
Metal Li Na K Rb Cs
Sodium Na 1s22s22p63s1
Potassium K 1s22s22p63s23p64s1 Colour Crimson Yellow Violet Red Blue
Rubidium Rb 1s22s22p63s23p63d104s24p65s1 red violet
Caesium Cs 1s22s22p63s23p63d104s2 λ/nm 670.8 589.2 766.5 780.0 455.5
6 10 2 6 1 1
4p 4d 5s 5p 6s or [Xe] 6s Alkali metals can therefore, be detected by
Francium Fr [Rn]7s 1 the respective flame tests and can be
determined by flame photometry or atomic
10.1.2 Atomic and Ionic Radii absorption spectroscopy. These elements when
The alkali metal atoms have the largest sizes irradiated with light, the light energy absorbed
in a particular period of the periodic table. With may be sufficient to make an atom lose electron.
THE s-BLOCK ELEMENTS 293
(ii) Reactivity towards water: The alkali the highest hydration enthalpy which
0
metals react with water to form hydroxide accounts for its high negative E value and
and dihydrogen. its high reducing power.
2 M + 2H2O → 2 M+ + 2OH− + H2
(M = an alkali metal) Problem 10.2
0 – –
It may be noted that although lithium has The E for Cl 2/Cl is +1.36, for I2/I is
0 + +
most negative E value (Table 10.1), its + 0.53, for Ag /Ag is +0.79, Na /Na is
+
reaction with water is less vigorous than –2.71 and for Li /Li is – 3.04. Arrange
that of sodium which has the least negative the following ionic species in decreasing
0
E value among the alkali metals. This order of reducing strength:
behaviour of lithium is attributed to its – –
I , Ag, Cl , Li, Na
small size and very high hydration energy.
Other metals of the group react explosively Solution
– –
with water. The order is Li > Na > I > Ag > Cl
They also react with proton donors such
as alcohol, gaseous ammonia and alkynes. (vi) Solutions in liquid ammonia: The alkali
(iii) Reactivity towards dihydrogen: The metals dissolve in liquid ammonia giving
alkali metals react with dihydrogen at deep blue solutions which are conducting
about 673K (lithium at 1073K) to form in nature.
hydrides. All the alkali metal hydrides are M + (x + y)NH3 →[M(NH3 )x ]+ + [e(NH3 )y ]−
ionic solids with high melting points. The blue colour of the solution is due to
the ammoniated electron which absorbs
2 M + H2 → 2 M + H −
energy in the visible region of light and thus
(iv) Reactivity towards halogens : The alkali imparts blue colour to the solution. The
metals readily react vigorously with solutions are paramagnetic and on
+ –
halogens to form ionic halides, M X . standing slowly liberate hydrogen resulting
However, lithium halides are somewhat in the formation of amide.
covalent. It is because of the high
polarisation capability of lithium ion (The M + (am ) + e − + NH3 (1) → MNH2(am ) + ½H2 (g)
distortion of electron cloud of the anion by (where ‘am’ denotes solution in ammonia.)
+
the cation is called polarisation). The Li ion In concentrated solution, the blue colour
is very small in size and has high tendency changes to bronze colour and becomes
to distort electron cloud around the diamagnetic.
negative halide ion. Since anion with large
size can be easily distorted, among halides, 10.1.7 Uses
lithium iodide is the most covalent in Lithium metal is used to make useful alloys,
nature. for example with lead to make ‘white metal’
(v) Reducing nature: The alkali metals are bearings for motor engines, with aluminium
strong reducing agents, lithium being the to make aircraft parts, and with magnesium
most and sodium the least powerful to make armour plates. It is used in
(Table 10.1). The standard electrode thermonuclear reactions. Lithium is also used
0
potential (E ) which measures the reducing to make electrochemical cells. Sodium is used
power represents the overall change : to make a Na/Pb alloy needed to make PbEt4
M(s) → M(g) sublimationenthalpy and PbMe4. These organolead compounds were
+
M(g) → M (g) + e −
ionizationenthalpy earlier used as anti-knock additives to petrol,
+ +
but nowadays vehicles use lead-free petrol.
M (g) + H2O → M (aq) hydrationenthalpy Liquid sodium metal is used as a coolant in
With the small size of its ion, lithium has fast breeder nuclear reactors. Potassium has
THE s-BLOCK ELEMENTS 295
a vital role in biological systems. Potassium The hydroxides which are obtained by the
chloride is used as a fertilizer. Potassium reaction of the oxides with water are all white
hydroxide is used in the manufacture of soft crystalline solids. The alkali metal hydroxides
soap. It is also used as an excellent absorbent are the strongest of all bases and dissolve freely
of carbon dioxide. Caesium is used in devising in water with evolution of much heat on
photoelectric cells. account of intense hydration.
Uses: It is used in (i) the manufacture of soap, found on the opposite sides of cell membranes.
paper, artificial silk and a number of chemicals, As a typical example, in blood plasma, sodium
–1
(ii) in petroleum refining, (iii) in the purification is present to the extent of 143 mmolL ,
of bauxite, (iv) in the textile industries for whereas the potassium level is only
–1
mercerising cotton fabrics, (v) for the 5 mmolL within the red blood cells. These
–1 +
preparation of pure fats and oils, and (vi) as a concentrations change to 10 mmolL (Na ) and
–1 +
laboratory reagent. 105 mmolL (K ). These ionic gradients
Sodium Hydrogencarbonate (Baking demonstrate that a discriminatory mechanism,
Soda), NaHCO3 called the sodium-potassium pump, operates
across the cell membranes which consumes
Sodium hydrogencarbonate is known as
more than one-third of the ATP used by a
baking soda because it decomposes on heating
resting animal and about 15 kg per 24 h in a
to generate bubbles of carbon dioxide (leaving
resting human.
holes in cakes or pastries and making them
light and fluffy). 10.6 GROUP 2 ELEMENTS : ALKALINE
Sodium hydrogencarbonate is made by EARTH METALS
saturating a solution of sodium carbonate with The group 2 elements comprise beryllium,
carbon dioxide. The white crystalline powder magnesium, calcium, strontium, barium and
of sodium hydrogencarbonate, being less radium. They follow alkali metals in the
soluble, gets separated out. periodic table. These (except beryllium) are
known as alkaline earth metals. The first
Na 2 CO3 + H2 O + CO2 → 2 NaHCO3
element beryllium differs from the rest of the
Sodium hydrogencarbonate is a mild members and shows diagonal relationship to
antiseptic for skin infections. It is used in fire aluminium. The atomic and physical
extinguishers. properties of the alkaline earth metals are
shown in Table 10.2.
10.5 BIOLOGICAL IMPORTANCE OF
SODIUM AND POTASSIUM 10.6.1 Electronic Configuration
A typical 70 kg man contains about 90 g of Na These elements have two electrons in the
and 170 g of K compared with only 5 g of iron s -orbital of the valence shell (Table 10.2). Their
and 0.06 g of copper. general electronic configuration may be
2
Sodium ions are found primarily on the represented as [noble gas] ns . Like alkali
outside of cells, being located in blood plasma metals, the compounds of these elements are
and in the interstitial fluid which surrounds also predominantly ionic.
the cells. These ions participate in the
Element Symbol Electronic
transmission of nerve signals, in regulating the configuration
flow of water across cell membranes and in the
transport of sugars and amino acids into cells. Beryllium Be 1s22s2
Sodium and potassium, although so similar Magnesium Mg 1s22s22p63s2
chemically, differ quantitatively in their ability Calcium Ca 1s22s22p63s23p64s2
to penetrate cell membranes, in their transport Strontium Sr 1s22s22p63s23p63d10
mechanisms and in their efficiency to activate 4s24p65s2
enzymes. Thus, potassium ions are the most Barium Ba 1s22s22p63s23p63d104s2
abundant cations within cell fluids, where they 4p64d105s25p66s2 or
activate many enzymes, participate in the [Xe]6s2
oxidation of glucose to produce ATP and, with Radium Ra [Rn]7s2
sodium, are responsible for the transmission
of nerve signals. 10.6.2 Atomic and Ionic Radii
There is a very considerable variation in the The atomic and ionic radii of the alkaline earth
concentration of sodium and potassium ions metals are smaller than those of the
THE s-BLOCK ELEMENTS 299
Table 10.2 Atomic and Physical Properties of the Alkaline Earth Metals
corresponding alkali metals in the same increase in ionic size down the group.
periods. This is due to the increased nuclear 2+ 2+ 2+
Be > Mg > Ca > Sr > Ba
2+ 2+
strontium and barium impart characteristic (iv) Reactivity towards acids: The alkaline
brick red, crimson and apple green colours earth metals readily react with acids liberating
respectively to the flame. In flame the electrons dihydrogen.
are excited to higher energy levels and when M + 2HCl → MCl2 + H2
they drop back to the ground state, energy is
(v) Reducing nature: Like alkali metals, the
emitted in the form of visible light. The
alkaline earth metals are strong reducing
electrons in beryllium and magnesium are too
strongly bound to get excited by flame. Hence, agents. This is indicated by large negative
these elements do not impart any colour to the values of their reduction potentials
flame. The flame test for Ca, Sr and Ba is (Table 10.2). However their reducing power is
helpful in their detection in qualitative analysis less than those of their corresponding alkali
and estimation by flame photometry. The metals. Beryllium has less negative value
alkaline earth metals like those of alkali metals compared to other alkaline earth metals.
have high electrical and thermal conductivities However, its reducing nature is due to large
which are typical characteristics of metals. hydration energy associated with the small
2+
size of Be ion and relatively large value of the
10.6.6 Chemical Properties
atomization enthalpy of the metal.
The alkaline earth metals are less reactive than
(vi) Solutions in liquid ammonia: Like
the alkali metals. The reactivity of these
alkali metals, the alkaline earth metals dissolve
elements increases on going down the group.
in liquid ammonia to give deep blue black
(i) Reactivity towards air and water: solutions forming ammoniated ions.
Beryllium and magnesium are kinetically inert
2+
M + ( x + y ) NH3 → ⎡⎣M ( NH3 ) X ⎤⎦ + 2 ⎡⎣e ( NH3 ) Y ⎤⎦
–
to oxygen and water because of the formation
of an oxide film on their surface. However,
powdered beryllium burns brilliantly on From these solutions, the ammoniates,
2+
ignition in air to give BeO and Be 3 N 2 . [M(NH3)6] can be recovered.
Magnesium is more electropositive and burns
with dazzling brilliance in air to give MgO and 10.6.7 Uses
Mg3N2. Calcium, strontium and barium are Beryllium is used in the manufacture of alloys.
readily attacked by air to form the oxide and Copper -beryllium alloys are used in the
nitride. They also react with water with preparation of high strength springs. Metallic
increasing vigour even in cold to form beryllium is used for making windows of
hydroxides. X-ray tubes. Magnesium forms alloys with
(ii) Reactivity towards the halogens: All aluminium, zinc, manganese and tin.
the alkaline earth metals combine with halogen Magnesium-aluminium alloys being light in
at elevated temperatures forming their halides. mass are used in air -craft construction.
M + X 2 → MX 2 ( X = F, Cl, Br, l ) Magnesium (powder and ribbon) is used in
Thermal decomposition of (NH4)2BeF4 is the flash powders and bulbs, incendiary bombs
best route for the preparation of BeF2, and and signals. A suspension of magnesium
BeCl2 is conveniently made from the oxide. hydroxide in water (called milk of magnesia)
600 − 800K is used as antacid in medicine. Magnesium
BeO + C + Cl 2 BeCl 2 + CO carbonate is an ingredient of toothpaste.
(iii) Reactivity towards hydrogen: All the Calcium is used in the extraction of metals from
elements except beryllium combine with oxides which are difficult to reduce with
hydrogen upon heating to form their hydrides, carbon. Calcium and barium metals, owing
MH2. to their reactivity with oxygen and nitrogen at
BeH2, however, can be prepared by the reaction elevated temperatures, have often been used
of BeCl2 with LiAlH4. to remove air from vacuum tubes. Radium
salts are used in radiotherapy, for example, in
2BeCl 2 + LiAlH 4 → 2BeH 2 + LiCl + AlCl 3
the treatment of cancer.
THE s-BLOCK ELEMENTS 301
( )
heated to 393 K.
2Ca ( OH )2 + 2Cl2 → CaCl 2 + Ca OCl + 2H2O
Bleaching powder
2
2 ( CaSO4 .2H2O ) → 2 ( CaSO4 ) .H2 O + 3H2O
Uses: Above 393 K, no water of crystallisation is left
(i) It is used in the preparation of mortar, a and anhydrous calcium sulphate, CaSO4 is
building material. formed. This is known as ‘dead burnt plaster’.
304 CHEMISTRY
It has a remarkable property of setting with silicate (Ca 3 SiO 5 ) 51% and tricalcium
water. On mixing with an adequate quantity aluminate (Ca3Al2O6) 11%.
of water it forms a plastic mass that gets into a Setting of Cement: When mixed with water,
hard solid in 5 to 15 minutes. the setting of cement takes place to give a hard
Uses: mass. This is due to the hydration of the
The largest use of Plaster of Paris is in the molecules of the constituents and their
building industry as well as plasters. It is used rearrangement. The purpose of adding
for immoblising the affected part of organ where gypsum is only to slow down the process of
there is a bone fracture or sprain. It is also setting of the cement so that it gets sufficiently
employed in dentistry, in ornamental work and hardened.
for making casts of statues and busts. Uses: Cement has become a commodity of
Cement: Cement is an important building national necessity for any country next to iron
material. It was first introduced in England in and steel. It is used in concrete and reinforced
1824 by Joseph Aspdin. It is also called concrete, in plastering and in the construction
Portland cement because it resembles with the of bridges, dams and buildings.
natural limestone quarried in the Isle of
10.10 BIOLOGICAL IMPORTANCE OF
Portland, England.
MAGNESIUM AND CALCIUM
Cement is a product obtained by An adult body contains about 25 g of Mg and
combining a material rich in lime, CaO with 1200 g of Ca compared with only 5 g of iron
other material such as clay which contains and 0.06 g of copper. The daily requirement
silica, SiO 2 along with the oxides of in the human body has been estimated to be
aluminium, iron and magnesium. The average 200 – 300 mg.
composition of Portland cement is : CaO, 50-
All enzymes that utilise ATP in phosphate
60%; SiO2, 20-25%; Al2O3, 5-10%; MgO, 2-
transfer require magnesium as the cofactor.
3%; Fe2O3, 1-2% and SO3, 1-2%. For a good
The main pigment for the absorption of light
quality cement, the ratio of silica (SiO2) to
in plants is chlorophyll which contains
alumina (Al2O3) should be between 2.5 and 4
magnesium. About 99 % of body calcium is
and the ratio of lime (CaO) to the total of the
present in bones and teeth. It also plays
oxides of silicon (SiO2) aluminium (Al2O3) and
important roles in neuromuscular function,
iron (Fe2O3) should be as close as possible to
interneuronal transmission, cell membrane
2.
integrity and blood coagulation. The calcium
The raw materials for the manufacture of concentration in plasma is regulated at about
cement are limestone and clay. When clay and 100 mgL–1. It is maintained by two hormones:
lime are strongly heated together they fuse and calcitonin and parathyroid hormone. Do you
react to form ‘cement clinker’. This clinker is know that bone is not an inert and unchanging
mixed with 2-3% by weight of gypsum substance but is continuously being
(CaSO4·2H2O) to form cement. Thus important solubilised and redeposited to the extent of
ingredients present in Portland cement are 400 mg per day in man? All this calcium
dicalcium silicate (Ca2SiO4) 26%, tricalcium passes through the plasma.
SUMMARY
The s-Block of the periodic table constitutes Group1 (alkali metals) and Group 2
(alkaline earth metals). They are so called because their oxides and hydroxides are alkaline
in nature. The alkali metals are characterised by one s-electron and the alkaline earth
metals by two s-electrons in the+ valence shell of their atoms. These are highly reactive
2+
metals forming monopositive (M ) and dipositve (M ) ions respectively.
THE s-BLOCK ELEMENTS 305
There is a regular trend in the physical and chemical properties of the alkali metal
with increasing atomic numbers. The atomic and ionic sizes increase and the ionization
enthalpies decrease systematically down the group. Somewhat similar trends are
observed among the properties of the alkaline earth metals.
The first element in each of these groups, lithium in Group 1 and beryllium in
Group 2 shows similarities in properties to the second member of the next group. Such
similarities are termed as the ‘diagonal relationship’ in the periodic table. As such
these elements are anomalous as far as their group characteristics are concerned.
The alkali metals are silvery white, soft and low melting. They are highly reactive.
The compounds of alkali metals are predominantly ionic. Their oxides and hydroxides
are soluble in water forming strong alkalies. Important compounds of sodium includes
sodium carbonate, sodium chloride, sodium hydroxide and sodium hydrogencarbonate.
Sodium hydroxide is manufactured by Castner-Kellner process and sodium carbonate
by Solvay process.
The chemistry of alkaline earth metals is very much like that of the alkali metals.
However, some differences arise because of reduced atomic and ionic sizes and increased
cationic charges in case of alkaline earth metals. Their oxides and hydroxides are less
basic than the alkali metal oxides and hydroxides. Industrially important compounds of
calcium include calcium oxide (lime), calcium hydroxide (slaked lime), calcium sulphate
(Plaster of Paris), calcium carbonate (limestone) and cement. Portland cement is an
important constructional material. It is manufactured by heating a pulverised mixture
of limestone and clay in a rotary kiln. The clinker thus obtained is mixed with some
gypsum (2-3%) to give a fine powder of cement. All these substances find variety of uses
in different areas.
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.
EXERCISES
10.1 What are the common physical and chemical features of alkali metals ?
10.2 Discuss the general characteristics and gradation in properties of alkaline earth
metals.
10.3 Why are alkali metals not found in nature ?
10.4 Find out the oxidation state of sodium in Na2O2.
10.5 Explain why is sodium less reactive than potassium.
10.6 Compare the alkali metals and alkaline earth metals with respect to (i) ionisation
enthalpy (ii) basicity of oxides and (iii) solubility of hydroxides.
10.7 In what ways lithium shows similarities to magnesium in its chemical behaviour?
10.8 Explain why can alkali and alkaline earth metals not be obtained by chemical
reduction methods?
10.9 Why are potassium and caesium, rather than lithium used in photoelectric cells?
10.10 When an alkali metal dissolves in liquid ammonia the solution can acquire
different colours. Explain the reasons for this type of colour change.
10.11 Beryllium and magnesium do not give colour to flame whereas other alkaline
earth metals do so. Why ?
10.12 Discuss the various reactions that occur in the Solvay process.
10.13 Potassium carbonate cannot be prepared by Solvay process. Why ?
10.14 Why is Li2CO3 decomposed at a lower temperature whereas Na2CO3 at higher
temperature?
306 CHEMISTRY
10.15 Compare the solubility and thermal stability of the following compounds of the
alkali metals with those of the alkaline earth metals. (a) Nitrates (b) Carbonates
(c) Sulphates.
10.16 Starting with sodium chloride how would you proceed to prepare (i) sodium metal
(ii) sodium hydroxide (iii) sodium peroxide (iv) sodium carbonate ?
10.17 What happens when (i) magnesium is burnt in air (ii) quick lime is heated with
silica (iii) chlorine reacts with slaked lime (iv) calcium nitrate is heated ?
10.18 Describe two important uses of each of the following : (i) caustic soda (ii) sodium
carbonate (iii) quicklime.
10.19 Draw the structure of (i) BeCl2 (vapour) (ii) BeCl2 (solid).
10.20 The hydroxides and carbonates of sodium and potassium are easily soluble in
water while the corresponding salts of magnesium and calcium are sparingly
soluble in water. Explain.
10.21 Describe the importance of the following : (i) limestone (ii) cement (iii) plaster of
paris.
10.22 Why are lithium salts commonly hydrated and those of the other alkali ions
usually anhydrous?
10.23 Why is LiF almost insoluble in water whereas LiCl soluble not only in water but
also in acetone ?
10.24 Explain the significance of sodium, potassium, magnesium and calcium in
biological fluids.
10.25 What happens when
(i) sodium metal is dropped in water ?
(ii) sodium metal is heated in free supply of air ?
(iii) sodium peroxide dissolves in water ?
10.26 Comment on each of the following observations:
+ + +
(a) The mobilities of the alkali metal ions in aqueous solution are Li < Na < K
+ +
< Rb < Cs
(b) Lithium is the only alkali metal to form a nitride directly.
0 2+ –
(c) E for M (aq) + 2e → M(s) (where M = Ca, Sr or Ba) is nearly constant.
10.27 State as to why
(a) a solution of Na2CO3 is alkaline ?
(b) alkali metals are prepared by electrolysis of their fused chlorides ?
(c) sodium is found to be more useful than potassium ?
10.28 Write balanced equations for reactions between
(a) Na2O2 and water
(b) KO2 and water
(c) Na2O and CO2.
10.29 How would you explain the following observations?
(i) BeO is almost insoluble but BeSO4 in soluble in water,
(ii) BaO is soluble but BaSO4 is insoluble in water,
(iii) LiI is more soluble than KI in ethanol.
10.30 Which of the alkali metal is having least melting point ?
(a) Na (b) K (c) Rb (d) Cs
10.31 Which one of the following alkali metals gives hydrated salts ?
(a) Li (b) Na (c) K (d) Cs
10.32 Which one of the alkaline earth metal carbonates is thermally the most stable ?
(a) MgCO3 (b) CaCO3 (c) SrCO3 (d) BaCO3
THE p-BLOCK ELEMENTS 307
UNIT 11
Table 11.1 General Electronic Configuration and Oxidation States of p-Block Elements
Group 13 14 15 16 17 18
General
electronic ns2np1 ns2np2 ns2np3 ns2np4 ns2np5 ns2np6
configuration (1s2 for He)
First member
of the B C N O F He
group
Group
oxidation +3 +4 +5 +6 +7 +8
state
Other
oxidation +1 +2, – 4 +3, – 3 +4, +2, –2 +5, + 3, +1, –1 +6, +4, +2
states
The relative stabilities of these two oxidation The first member of p-block differs from the
states – group oxidation state and two unit less remaining members of their corresponding
than the group oxidation state – may vary from group in two major respects. First is the size
group to group and will be discussed at and all other properties which depend on size.
appropriate places. Thus, the lightest p-block elements show the
It is interesting to note that the non-metals same kind of differences as the lightest s-block
and metalloids exist only in the p-block of the elements, lithium and beryllium. The second
periodic table. The non-metallic character of important difference, which applies only to the
elements decreases down the group. In fact the p-block elements, arises from the effect of d-
heaviest element in each p-block group is the orbitals in the valence shell of heavier elements
most metallic in nature. This change from non- (starting from the third period onwards) and
metallic to metallic character brings diversity their lack in second period elements. The
in the chemistry of these elements depending second period elements of p-groups starting
from boron are restricted to a maximum
on the group to which they belong.
covalence of four (using 2s and three 2p
In general, non-metals have higher ionisation orbitals). In contrast, the third period elements
enthalpies and higher electronegativities than of p-groups with the electronic configuration
the metals. Hence, in contrast to metals which n
3s23p have the vacant 3d orbitals lying
readily form cations, non-metals readily form between the 3p and the 4s levels of energy.
anions. The compounds formed by highly Using these d-orbitals the third period
reactive non-metals with highly reactive metals elements can expand their covalence above
are generally ionic because of large differences four. For example, while boron forms only
in their electronegativities. On the other hand, – 3–
[BF 4] , aluminium gives [AlF 6] ion. The
compounds formed between non-metals presence of these d-orbitals influences the
themselves are largely covalent in character chemistry of the heavier elements in a number
because of small differences in their of other ways. The combined effect of size and
electronegativities. The change of non-metallic availability of d orbitals considerably
to metallic character can be best illustrated by influences the ability of these elements to form
the nature of oxides they form. The non-metal π bonds. The first member of a group differs
oxides are acidic or neutral whereas metal from the heavier members in its ability to form
oxides are basic in nature. pπ - pπ multiple bonds to itself ( e.g., C=C, C≡C,
THE p-BLOCK ELEMENTS 309
N≡N) and to other second row elements (e.g., 11.1.1 Electronic Configuration
C=O, C=N, C≡N, N=O). This type of π - bonding The outer electronic configuration of these
2 1
is not particularly strong for the heavier elements is ns np . A close look at the
p-block elements. The heavier elements do form electronic configuration suggests that while
π bonds but this involves d orbitals (dπ – pπ boron and aluminium have noble gas
or dπ –dπ ). As the d orbitals are of higher core, gallium and indium have noble gas plus
energy than the p orbitals, they contribute less 10 d-electrons, and thallium has noble gas
to the overall stability of molecules than does plus 14 f- electrons plus 10 d-electrons cores.
pπ - pπ bonding of the second row elements. Thus, the electronic structures of these
However, the coordination number in species elements are more complex than for the first
of heavier elements may be higher than for two groups of elements discussed in unit 10.
the first element in the same oxidation state. This difference in electronic structures affects
For example, in +5 oxidation state both N and the other properties and consequently the
–
P form oxoanions : NO3 (three-coordination chemistry of all the elements of this group.
with π – bond involving one nitrogen p-orbital) 11.1.2 Atomic Radii
and PO34− (four-coordination involving s, p and On moving down the group, for each successive
d orbitals contributing to the π – bond). In member one extra shell of electrons is added
this unit we will study the chemistry of group and, therefore, atomic radius is expected to
13 and 14 elements of the periodic table. increase. However, a deviation can be seen.
Atomic radius of Ga is less than that of Al. This
11.1 GROUP 13 ELEMENTS: THE BORON
can be understood from the variation in the
FAMILY
inner core of the electronic configuration. The
This group elements show a wide variation in presence of additional 10 d-electrons offer
properties. Boron is a typical non-metal, only poor screening effect (Unit 2) for the outer
aluminium is a metal but shows many electrons from the increased nuclear charge in
chemical similarities to boron, and gallium, gallium. Consequently, the atomic radius of
indium and thallium are almost exclusively gallium (135 pm) is less than that of
metallic in character. aluminium (143 pm).
Boron is a fairly rare element, mainly 11.1.3 Ionization Enthalpy
occurs as orthoboric acid, (H3BO3), borax,
The ionisation enthalpy values as expected
Na2B4O7·10H2O, and kernite, Na2B4O7·4H2O.
from the general trends do not decrease
In India borax occurs in Puga Valley (Ladakh)
smoothly down the group. The decrease from
and Sambhar Lake (Rajasthan). The
B to Al is associated with increase in size. The
abundance of boron in earth crust is less than
observed discontinuity in the ionisation
0.0001% by mass. There are two isotopic
10 11 enthalpy values between Al and Ga, and
forms of boron B (19%) and B (81%).
between In and Tl are due to inability of d- and
Aluminium is the most abundant metal and
f-electrons ,which have low screening effect, to
the third most abundant element in the earth’s compensate the increase in nuclear charge.
crust (8.3% by mass) after oxygen (45.5%) and
The order of ionisation enthalpies, as
Si (27.7%). Bauxite, Al2O3. 2H2O and cryolite,
Na 3 AlF 6 are the important minerals of expected, is Δi H1<Δi H2<Δi H3. The sum of the
aluminium. In India it is found as mica in first three ionisation enthalpies for each of the
Madhya Pradesh, Karnataka, Orissa and elements is very high. Effect of this will be
Jammu. Gallium, indium and thallium are less apparent when you study their chemical
abundant elements in nature. properties.
The atomic, physical and chemical 11.1.4 Electronegativity
properties of these elements are discussed Down the group, electronegativity first
below. decreases from B to Al and then increases
310 CHEMISTRY
marginally (Table 11.2). This is because of the only covalent compounds. But as we move from
discrepancies in atomic size of the elements. B to Al, the sum of the first three ionisation
enthalpies of Al considerably decreases, and
11.1.5 Physical Properties 3+
is therefore able to form Al ions. In fact,
Boron is non-metallic in nature. It is extremely aluminium is a highly electropositive metal.
hard and black coloured solid. It exists in many However, down the group, due to poor
allotropic forms. Due to very strong crystalline shielding effect of intervening d and f orbitals,
lattice, boron has unusually high melting point. the increased effective nuclear charge holds ns
Rest of the members are soft metals with low electrons tightly (responsible for inter pair
melting point and high electrical conductivity. effect) and thereby, restricting their
It is worthwhile to note that gallium with participation in bonding. As a result of this,
unusually low melting point (303 K), could only p-orbital electron may be involved in
exist in liquid state during summer. Its high bonding. In fact in Ga, In and Tl, both +1 and
boiling point (2676 K) makes it a useful +3 oxidation states are observed. The relative
material for measuring high temperatures. stability of +1 oxidation state progressively
Density of the elements increases down the increases for heavier elements: Al<Ga<In<Tl. In
group from boron to thallium. thallium +1 oxidation state is predominant
whereas the +3 oxidation state is highly
11.1.6 Chemical Properties oxidising in character. The compounds in
Oxidation state and trends in chemical +1 oxidation state, as expected from energy
reactivity considerations, are more ionic than those in
Due to small size of boron, the sum of its first +3 oxidation state.
three ionization enthalpies is very high. This In trivalent state, the number of electrons
prevents it to form +3 ions and forces it to form around the central atom in a molecule
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
2 1 2 1 10 2 1 10 2 1
Electronic [He]2s 2p [Ne]3s 3p [Ar]3d 4s 4p [Kr]4d 5s 5p [Xe]4f145d106s26p1
Atomic radius/pma (85) 143 135 167 170
Ionic radius (27) 53.5 62.0 80.0 88.5
M3+/pmb
Ionic radius - - 120 140 150
M+/pm
Ionization Δi H 1 801 577 579 558 589
enthalpy Δi H 2 2427 1816 1979 1820 1971
(kJ mol–1) Δi H 3 3659 2744 2962 2704 2877
Electronegativity c 2.0 1.5 1.6 1.7 1.8
–3
Density /g cm 2.35 2.70 5.90 7.31 11.85
at 298 K
Melting point / K 2453 933 303 430 576
Boiling point / K 3923 2740 2676 2353 1730
V 3+
E / V for (M /M) - –1.66 –0.56 –0.34 +1.26
V +
E / V for (M /M) - +0.55 -0.79(acid) –0.18 –0.34
–1.39(alkali)
a b c
Metallic radius, 6-coordination, Pauling scale,
THE p-BLOCK ELEMENTS 311
2E ( s ) + 3O2 ( g ) ⎯⎯→
Δ
2E 2 O3 ( s )
AlCl3 achieves stability by forming a dimer
2E ( s ) + N 2 ( g ) ⎯⎯→
Δ
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.
In trivalent state most of the compounds
(ii) Reactivity towards acids and alkalies
being covalent are hydrolysed in water. For
Boron does not react with acids and alkalies
example, the trichlorides on hyrolysis in water
− even at moderate temperature; but aluminium
form tetrahedral ⎡⎣ M ( OH )4 ⎤⎦ species; the dissolves in mineral acids and aqueous alkalies
3
hybridisation state of element M is sp . and thus shows amphoteric character.
Aluminium chloride in acidified aqueous Aluminium dissolves in dilute HCl and
3+
solution forms octahedral ⎡⎣ Al ( H2 O )6 ⎤⎦ ion. liberates dihydrogen.
3+ –
In this complex ion, the 3d orbitals of Al are 2Al(s) + 6HCl (aq) → 2Al (aq) + 6Cl (aq)
involved and the hybridisation state of Al is + 3H2 (g)
sp3d2. However, concentrated nitric acid renders
aluminium passive by forming a protective
Problem 11.1 oxide layer on the surface.
V
Standard electrode potential values, E Aluminium also reacts with aqueous alkali
3+ 3+
for Al /Al is –1.66 V and that of Tl /Tl and liberates dihydrogen.
is +1.26 V. Predict about the formation of
3+
M ion in solution and compare the 2Al (s) + 2NaOH(aq) + 6H2O(l)
electropositive character of the two ↓
+ –
metals. 2 Na [Al(OH)4] (aq) + 3H2(g)
Sodium
Solution tetrahydroxoaluminate(III)
Standard electrode potential values for two (iii) Reactivity towards halogens
half cell reactions suggest that aluminium These elements react with halogens to form
3+
has high tendency to make Al (aq) ions, trihalides (except Tl I3).
3+
whereas Tl is not only unstable in
2E(s) + 3 X2 (g) → 2EX3 (s) (X = F, Cl, Br, I)
312 CHEMISTRY
Germanium exists only in traces. Tin occurs due to the presence of completely filled d and f
mainly as cassiterite, SnO 2 and lead as orbitals in heavier members.
galena, PbS.
11.5.3 Ionization Enthalpy
Ultrapure form of germanium and silicon
The first ionization enthalpy of group 14
are used to make transistors and
members is higher than the corresponding
semiconductor devices.
members of group 13. The influence of inner
The important atomic and physical core electrons is visible here also. In general the
properties of the group 14 elements along ionisation enthalpy decreases down the group.
with their electronic configuration are given Small decrease in ΔiH from Si to Ge to Sn and
in Table 11.2 Some of the atomic, physical slight increase in Δi H from Sn to Pb is the
and chemical properties are discussed consequence of poor shielding effect of
below: intervening d and f orbitals and increase in size
11.5.1 Electronic Configuration of the atom.
The valence shell electronic configuration of 11.5.4 Electronegativity
2 2
these elements is ns np . The inner core of the Due to small size, the elements of this group
electronic configuration of elements in this are slightly more electronegative than group
group also differs. 13 elements. The electronegativity values for
11.5.2 Covalent Radius elements from Si to Pb are almost the same.
There is a considerable increase in covalent 11.5.5 Physical Properties
radius from C to Si, thereafter from Si to Pb a All group 14 members are solids. Carbon and
small increase in radius is observed. This is silicon are non-metals, germanium is a metalloid,
Table 11.3 Atomic and Physical Properties of Group 14 Elements
Element
Property Carbon Silicon Germanium Tin Lead
C Si Ge Sn Pb
Atomic Number 6 14 32 50 82
–1
Atomic mass (g mol ) 12.01 28.09 72.60 118.71 207.2
2 2 2 2 10 2 2 10 2 2 14 2 2
Electronic [He]2s 2p [Ne]3s 3p [Ar]3d 4s 4p [Kr]4d 5s 5p [Xe]4f 5d6s 6p
configuration
a
Covalent radius/pm 77 118 122 140 146
4+ b
Ionic radius M /pm – 40 53 69 78
2+ b
Ionic radius M /pm – – 73 118 119
Ionization Δi H 1 1086 786 761 708 715
enthalpy/ Δi H 2 2352 1577 1537 1411 1450
kJ mol–1 Δi H 3 4620 3228 3300 2942 3081
Δi H 4 6220 4354 4409 3929 4082
c
Electronegativity 2.5 1.8 1.8 1.8 1.9
d –3 e f
Density /g cm 3.51 2.34 5.32 7.26 11.34
Melting point/K 4373 1693 1218 505 600
Boiling point/K – 3550 3123 2896 2024
14 16 –5 –5
Electrical resistivity/ 10 –10 50 50 10 2 × 10
ohm cm (293 K)
a IV b c d e
for M oxidation state; 6–coordination; Pauling scale; 293 K; for diamond; for graphite, density is
f
2.22; β-form (stable at room temperature)
316 CHEMISTRY
whereas tin and lead are soft metals with low those in lower oxidation states. The dioxides
melting points. Melting points and boiling points — CO2, SiO2 and GeO2 are acidic, whereas
of group 14 elements are much higher than those SnO2 and PbO2 are amphoteric in nature.
of corresponding elements of group 13. Among monoxides, CO is neutral, GeO is
distinctly acidic whereas SnO and PbO are
11.5.6 Chemical Properties
amphoteric.
Oxidation states and trends in chemical
reactivity Problem 11.5
The group 14 elements have four electrons in Select the member(s) of group 14 that
outermost shell. The common oxidation states (i) forms the most acidic dioxide, (ii) is
exhibited by these elements are +4 and +2. commonly found in +2 oxidation state,
Carbon also exhibits negative oxidation states. (iii) used as semiconductor.
Since the sum of the first four ionization
enthalpies is very high, compounds in +4 Solution
oxidation state are generally covalent in nature. (i) carbon (ii) lead
In heavier members the tendency to show +2 (iii) silicon and germanium
oxidation state increases in the sequence
2
Ge<Sn<Pb. It is due to the inability of ns (ii) Reactivity towards water
electrons of valence shell to participate in
bonding. The relative stabilities of these two Carbon, silicon and germanium are not
oxidation states vary down the group. Carbon affected by water. Tin decomposes steam to
and silicon mostly show +4 oxidation state. form dioxide and dihydrogen gas.
Germanium forms stable compounds in +4 Δ
Sn + 2H2O ⎯
→ SnO2 + 2H2
state and only few compounds in +2 state. Tin
forms compounds in both oxidation states (Sn Lead is unaffected by water, probably
in +2 state is a reducing agent). Lead because of a protective oxide film formation.
compounds in +2 state are stable and in +4 (iii) Reactivity towards halogen
state are strong oxidising agents. In tetravalent
These elements can form halides of formula
state the number of electrons around the
MX2 and MX4 (where X = F, Cl, Br, I). Except
central atom in a molecule (e.g., carbon in CCl4)
is eight. Being electron precise molecules, they carbon, all other members react directly with
are normally not expected to act as electron halogen under suitable condition to make
acceptor or electron donor species. Although halides. Most of the MX4 are covalent in nature.
carbon cannot exceed its covalence more than The central metal atom in these halides
3
4, other elements of the group can do so. It is undergoes sp hybridisation and the molecule
because of the presence of d orbital in them. is tetrahedral in shape. Exceptions are SnF4
Due to this, their halides undergo hydrolysis and PbF4, which are ionic in nature. PbI4 does
and have tendency to form complexes by not exist because Pb—I bond initially formed
accepting electron pairs from donor species. For during the reaction does not release enough
2– 2– 2
example, the species like, SiF6 , [GeCl6] , energy to unpair 6s electrons and excite one
2–
[Sn(OH)6] exist where the hybridisation of the of them to higher orbital to have four unpaired
3 2
central atom is sp d . electrons around lead atom. Heavier members
(i) Reactivity towards oxygen Ge to Pb are able to make halides of formula
MX2. Stability of dihalides increases down the
All members when heated in oxygen form
group. Considering the thermal and chemical
oxides. There are mainly two types of oxides,
i.e., monoxide and dioxide of formula MO and stability, GeX4 is more stable than GeX2,
MO2 respectively. SiO only exists at high whereas PbX2 is more than PbX4. Except CCl4,
temperature. Oxides in higher oxidation states other tetrachlorides are easily hydrolysed
of elements are generally more acidic than by water because the central atom can
THE p-BLOCK ELEMENTS 317
accommodate the lone pair of electrons from Carbon also has unique ability to form
oxygen atom of water molecule in d orbital. pπ– pπ multiple bonds with itself and with other
Hydrolysis can be understood by taking atoms of small size and high electronegativity.
the example of SiCl4. It undergoes hydrolysis Few examples of multiple bonding are: C=C,
by initially accepting lone pair of electrons C ≡ C, C = O, C = S, and C ≡ N. Heavier elements
from water molecule in d orbitals of Si, finally do not form pπ– pπ bonds because their atomic
leading to the formation of Si(OH)4 as shown orbitals are too large and diffuse to have
below : effective overlapping.
Carbon atoms have the tendency to link
with one another through covalent bonds to
form chains and rings. This property is called
catenation. This is because C—C bonds are
very strong. Down the group the size increases
and electronegativity decreases, and, thereby,
tendency to show catenation decreases. This
can be clearly seen from bond enthalpies
values. The order of catenation is C > > Si >
Ge ≈ Sn. Lead does not show catenation.
–1
Bond Bond enthalpy / kJ mol
Problem 11. 6 C—C 348
2– 2–
[SiF6] is known whereas [SiCl6] not. Si —Si 297
Give possible reasons. Ge—Ge 260
Solution Sn—Sn 240
The main reasons are :
(i) six large chloride ions cannot be Due to property of catenation and pπ– pπ
accommodated around Si due to
4+ bond formation, carbon is able to show
limitation of its size. allotropic forms.
(ii) interaction between lone pair of 11.7 ALLOTROPES OF CARBON
4+
chloride ion and Si is not very strong.
Carbon exhibits many allotropic forms; both
crystalline as well as amorphous. Diamond
11.6 IMPORTANT TRENDS AND and graphite are two well-known crystalline
ANOMALOUS BEHAVIOUR OF forms of carbon. In 1985, third form of carbon
CARBON known as fullerenes was discovered by
Like first member of other groups, carbon H.W.Kroto, E.Smalley and R.F.Curl. For this
also differs from rest of the members of its discovery they were awarded the Nobel Prize
group. It is due to its smaller size, higher in 1996.
electronegativity, higher ionisation enthalpy
and unavailability of d orbitals. 11.7.1 Diamond
In carbon, only s and p orbitals are It has a crystalline lattice. In diamond each
3
available for bonding and, therefore, it can carbon atom undergoes sp hybridisation and
accommodate only four pairs of electrons linked to four other carbon atoms by using
around it. This would limit the maximum hybridised orbitals in tetrahedral fashion. The
covalence to four whereas other members can C–C bond length is 154 pm. The structure
expand their covalence due to the presence of extends in space and produces a rigid three-
d orbitals. dimensional network of carbon atoms. In this
318 CHEMISTRY
Fig. 11.3 The structure of diamond Fig 11.4 The structure of graphite
structure (Fig. 11.3) directional covalent bonds therefore, graphite conducts electricity along
are present throughout the lattice. the sheet. Graphite cleaves easily between the
It is very difficult to break extended covalent layers and, therefore, it is very soft and slippery.
bonding and, therefore, diamond is a hardest For this reason graphite is used as a dry
substance on the earth. It is used as an lubricant in machines running at high
abrasive for sharpening hard tools, in making temperature, where oil cannot be used as a
dies and in the manufacture of tungsten lubricant.
filaments for electric light bulbs. 11.7.3 Fullerenes
Problem 11.7 Fullerenes are made by the heating of graphite
Diamond is covalent, yet it has high in an electric arc in the presence of inert gases
melting point. Why ? such as helium or argon. The sooty material
n
formed by condensation of vapourised C small
Solution molecules consists of mainly C60 with smaller
Diamond has a three-dimensional quantity of C 70 and traces of fullerenes
network involving strong C—C bonds, consisting of even number of carbon atoms up
which are very difficult to break and, in to 350 or above. Fullerenes are the only pure
turn has high melting point. form of carbon because they have smooth
structure without having ‘dangling’ bonds.
11.7.2 Graphite Fullerenes are cage like molecules. C 60
Graphite has layered structure (Fig.11.4). molecule has a shape like soccer ball and
Layers are held by van der Waals forces and called Buckminsterfullerene (Fig. 11.5).
distance between two layers is 340 pm. Each It contains twenty six- membered rings and
layer is composed of planar hexagonal rings twelve five membered rings. A six membered
of carbon atoms. C—C bond length within the ring is fused with six or five membered rings
layer is 141.5 pm. Each carbon atom in but a five membered ring can only fuse with
2
hexagonal ring undergoes sp hybridisation six membered rings. All the carbon atoms are
2
and makes three sigma bonds with three equal and they undergo sp hybridisation.
neighbouring carbon atoms. Fourth electron Each carbon atom forms three sigma bonds
forms a π bond. The electrons are delocalised with other three carbon atoms. The remaining
over the whole sheet. Electrons are mobile and, electron at each carbon is delocalised in
THE p-BLOCK ELEMENTS 319
molecular orbitals, which in turn give aromatic filters to remove organic contaminators and in
character to molecule. This ball shaped airconditioning system to control odour.
molecule has 60 vertices and each one is Carbon black is used as black pigment in
occupied by one carbon atom and it also black ink and as filler in automobile tyres. Coke
contains both single and double bonds with is used as a fuel and largely as a reducing
C–C distances of 143.5 pm and 138.3 pm agent in metallurgy. Diamond is a precious
respectively. Spherical fullerenes are also called stone and used in jewellery. It is measured in
bucky balls in short. carats (1 carat = 200 mg).
11.8 SOME IMPORTANT COMPOUNDS OF
CARBON AND SILICON
Oxides of Carbon
Two important oxides of carbon are carbon
monoxide, CO and carbon dioxide, CO2.
11.8.1 Carbon Monoxide
Direct oxidation of C in limited supply of
oxygen or air yields carbon monoxide.
Δ
2C(s) + O2 (g) ⎯⎯⎯→ 2CO(g)
On small scale pure CO is prepared by
dehydration of formic acid with concentrated
H2SO4 at 373 K
373K
Fig.11.5 The structure of C 60, Buckminster- HCOOH ⎯⎯⎯⎯⎯
conc.H SO→ H2 O + CO
2 4
fullerene : Note that molecule has the
shape of a soccer ball (football). On commercial scale it is prepared by the
passage of steam over hot coke. The mixture
It is very important to know that graphite of CO and H2 thus produced is known as water
is thermodynamically most stable allotrope of gas or synthesis gas.
V
carbon and, therefore, Δf H of graphite is taken 473−1273K
V
as zero. Δf H values of diamond and fullerene, C ( s ) + H2 O ( g ) ⎯⎯⎯⎯⎯⎯⎯ → CO ( g ) + H2 ( g )
–1
C60 are 1.90 and 38.1 kJ mol , respectively. Water gas
Other forms of elemental carbon like carbon When air is used instead of steam, a mixture
black, coke, and charcoal are all impure forms of CO and N2 is produced, which is called
of graphite or fullerenes. Carbon black is producer gas.
obtained by burning hydrocarbons in a limited 1273K
supply of air. Charcoal and coke are obtained 2C(s) + O2 (g) + 4N 2 (g) ⎯⎯⎯⎯⎯ → 2CO(g)
by heating wood or coal respectively at high + 4N 2 (g)
temperatures in the absence of air. Producer gas
11.7.4 Uses of Carbon Water gas and producer gas are very
Graphite fibres embedded in plastic material important industrial fuels. Carbon monoxide
form high strength, lightweight composites. in water gas or producer gas can undergo
The composites are used in products such as further combustion forming carbon dioxide
tennis rackets, fishing rods, aircrafts and with the liberation of heat.
canoes. Being good conductor, graphite is used Carbon monoxide is a colourless,
for electrodes in batteries and industrial odourless and almost water insoluble gas. It
electrolysis. Crucibles made from graphite are is a powerful reducing agent and reduces
inert to dilute acids and alkalies. Being highly almost all metal oxides other than those of the
porous, activated charcoal is used in alkali and alkaline earth metals, aluminium
adsorbing poisonous gases; also used in water and a few transition metals. This property of
320 CHEMISTRY
CO is used in the extraction of many metals atmosphere, is removed from it by the process
from their oxides ores. known as photosynthesis. It is the process
Δ by which green plants convert atmospheric
Fe 2 O3 ( s ) + 3CO ( g ) ⎯⎯⎯→ 2Fe ( s ) + 3CO2 ( g )
CO2 into carbohydrates such as glucose. The
Δ
ZnO ( s ) + CO ( g ) ⎯⎯⎯→ Zn ( s ) + CO2 ( g ) overall chemical change can be expressed as:
In CO molecule, there are one sigma and hν
6CO2 +12H2 O ⎯⎯⎯⎯⎯⎯→ C6 H12 O6 + 6O2
two π bonds between carbon and oxygen, Chlorphyll
:C ≡ O: . Because of the presence of a lone pair + 6H2 O
on carbon, CO molecule acts as a donor and By this process plants make food for
reacts with certain metals when heated to form themselves as well as for animals and human
metal carbonyls. The highly poisonous beings. Unlike CO, it is not poisonous. But the
nature of CO arises because of its ability to increase in combustion of fossil fuels and
form a complex with haemoglobin, which decomposition of limestone for cement
is about 300 times more stable than the manufacture in recent years seem to increase
oxygen-haemoglobin complex. This prevents the CO2 content of the atmosphere. This may
haemoglobin in the red blood corpuscles from lead to increase in green house effect and
carrying oxygen round the body and ultimately thus, raise the temperature of the atmosphere
resulting in death. which might have serious consequences.
11.8.2 Carbon Dioxide Carbon dioxide can be obtained as a solid
in the form of dry ice by allowing the liquified
It is prepared by complete combustion of
CO2 to expand rapidly. Dry ice is used as a
carbon and carbon containing fuels in excess
refrigerant for ice-cream and frozen food.
of air.
Gaseous CO2 is extensively used to carbonate
Δ soft drinks. Being heavy and non-supporter
C(s) + O2 (g) ⎯⎯⎯→ CO2 (g)
of combustion it is used as fire extinguisher. A
Δ
CH4 (g) + 2O2 (g) ⎯⎯⎯→ CO2 (g) + 2H2 O(g) substantial amount of CO 2 is used to
manufacture urea.
In the laboratory it is conveniently
prepared by the action of dilute HCl on calcium In CO2 molecule carbon atom undergoes
carbonate. sp hybridisation. Two sp hybridised orbitals
of carbon atom overlap with two p orbitals of
CaCO3(s) + 2HCl (aq) → CaCl2 (aq) + CO2 (g) +
oxygen atoms to make two sigma bonds while
H2O(l)
other two electrons of carbon atom are involved
On commercial scale it is obtained by in pπ– pπ bonding with oxygen atom. This
heating limestone. results in its linear shape [with both C–O bonds
It is a colourless and odourless gas. Its low of equal length (115 pm)] with no dipole
solubility in water makes it of immense bio- moment. The resonance structures are shown
chemical and geo-chemical importance. With below:
water, it forms carbonic acid, H2CO3 which is
a weak dibasic acid and dissociates in two
steps: Resonance structures of carbon dioxide
– +
H2CO3(aq) + H2O(l) HCO3 (aq) + H3O (aq)
– 2– + 11.8.3 Silicon Dioxide, SiO2
HCO3 (aq) + H2O(l) CO3 (aq) + H3O (aq)
– 95% of the earth’s crust is made up of silica
H 2 CO 3/HCO 3 buffer system helps to and silicates. Silicon dioxide, commonly known
maintain pH of blood between 7.26 to 7.42. as silica, occurs in several crystallographic
Being acidic in nature, it combines with alkalies forms. Quartz, cristobalite and tridymite are
to form metal carbonates. some of the crystalline forms of silica, and they
Carbon dioxide, which is normally present are interconvertable at suitable temperature.
to the extent of ~ 0.03 % by volume in the Silicon dioxide is a covalent, three-dimensional
THE p-BLOCK ELEMENTS 321
network solid in which each silicon atom is substituted chlorosilane of formula MeSiCl3,
covalently bonded in a tetrahedral manner to Me2SiCl2, Me3SiCl with small amount of Me4Si
four oxygen atoms. Each oxygen atom in turn are formed. Hydrolysis of dimethyl-
covalently bonded to another silicon atoms as dichlorosilane, (CH 3 ) 2 SiCl 2 followed by
shown in diagram (Fig 11.6 ). Each corner is condensation polymerisation yields straight
shared with another tetrahedron. The entire chain polymers.
crystal may be considered as giant molecule
in which eight membered rings are formed with
alternate silicon and oxygen atoms.
Problem: 11.8
What are silicones ? (a) (b)
4–
Solution Fig. 11.7 (a) Tetrahedral structure of SiO 4
4–
anion; (b) Representation of SiO4 unit
Simple silicones consist of
neutralized by positively charged metal ions.
chains in which alkyl or phenyl groups If all the four corners are shared with other
occupy the remaining bonding positions tetrahedral units, three-dimensional network
on each silicon. They are hydrophobic is formed.
(water repellant) in nature. Two important man-made silicates are
glass and cement.
11.8.5 Silicates
11.8.6 Zeolites
A large number of silicates minerals exist in
nature. Some of the examples are feldspar, If aluminium atoms replace few silicon atoms
zeolites, mica and asbestos. The basic in three-dimensional network of silicon dioxide,
4– overall structure known as aluminosilicate,
structural unit of silicates is SiO4 (Fig.11.7)
in which silicon atom is bonded to four acquires a negative charge. Cations such as
+ +
oxygen atoms in tetrahedron fashion. In Na , K or Ca2+ balance the negative charge.
silicates either the discrete unit is present or Examples are feldspar and zeolites. Zeolites are
a number of such units are joined together widely used as a catalyst in petrochemical
via corners by sharing 1,2,3 or 4 oxygen industries for cracking of hydrocarbons and
atoms per silicate units. When silicate units isomerisation, e.g., ZSM-5 (A type of zeolite)
are linked together, they form chain, ring, used to convert alcohols directly into gasoline.
sheet or three-dimensional structures. Hydrated zeolites are used as ion exchangers
Negative charge on silicate structure is in softening of “hard” water.
SUMMARY
p-Block of the periodic table is unique in terms of having all types of elements – metals,
non-metals and metalloids. There are six groups of p-block elements in the periodic
2 1–6
table numbering from 13 to 18. Their valence shell electronic configuration is ns np
(except for He). Differences in the inner core of their electronic configuration greatly
influence their physical and chemical properties. As a consequence of this, a lot of
variation in properties among these elements is observed. In addition to the group oxidation
state, these elements show other oxidation states differing from the total number of valence
electrons by unit of two. While the group oxidation state is the most stable for the lighter
elements of the group, lower oxidation states become progressively more stable for the
heavier elements. The combined effect of size and availability of d orbitals considerably
THE p-BLOCK ELEMENTS 323
influences the ability of these elements to form π-bonds. While the lighter elements form
pπ –pπ bonds, the heavier ones form dπ–pπ or dπ –dπ bonds. Absence of d orbital in
second period elements limits their maximum covalence to 4 while heavier ones can
exceed this limit.
Boron is a typical non-metal and the other members are metals. The availability of 3
2 1
valence electrons (2s 2p ) for covalent bond formation using four orbitals (2s, 2px, 2py and
2pz) leads to the so called electron deficiency in boron compounds. This deficiency
makes them good electron acceptor and thus boron compounds behave as Lewis acids.
Boron forms covalent molecular compounds with dihydrogen as boranes, the simplest of
which is diborane, B2H6. Diborane contains two bridging hydrogen atoms between two
boron atoms; these bridge bonds are considered to be three-centre two-electron bonds.
The important compounds of boron with dioxygen are boric acid and borax. Boric acid,
B(OH)3 is a weak monobasic acid; it acts as a Lewis acid by accepting electrons from
hydroxyl ion. Borax is a white crystalline solid of formula Na2[B4O5(OH)4]·8H2O. The borax
bead test gives characteristic colours of transition metals.
Aluminium exhibits +3 oxidation state. With heavier elements +1 oxidation state gets
progressively stabilised on going down the group. This is a consequence of the so called
inert pair effect.
Carbon is a typical non-metal forming covalent bonds employing all its four valence
2 2
electrons (2s 2p ). It shows the property of catenation, the ability to form chains or
rings, not only with C–C single bonds but also with multiple bonds (C=C or C≡C). The
tendency to catenation decreases as C>>Si>Ge ~ Sn > Pb. Carbon provides one of the
best examples of allotropy. Three important allotropes of carbon are diamond, graphite
and fullerenes. The members of the carbon family mainly exhibit +4 and +2 oxidation
states; compouds in +4 oxidation states are generally covalent in nature. The tendency
to show +2 oxidation state increases among heavier elements. Lead in +2 state is stable
whereas in +4 oxidation state it is a strong oxidising agent. Carbon also exhibits negative
oxidation states. It forms two important oxides: CO and CO2. Carbon monoxide is neutral
whereas CO2 is acidic in nature. Carbon monoxide having lone pair of electrons on C
forms metal carbonyls. It is deadly poisonous due to higher stability of its haemoglobin
complex as compared to that of oxyhaemoglobin complex. Carbon dioxide as such is not
toxic. However, increased content of CO2 in atmosphere due to combustion of fossil fuels
and decomposition of limestone is feared to cause increase in ‘green house effect’. This,
in turn, raises the temperature of the atmosphere and causes serious complications.
Silica, silicates and silicones are important class of compounds and find applications
in industry and technology.
EXERCISES
11.9 What are electron deficient compounds ? Are BCl3 and SiCl 4 electron
deficient species ? Explain.
2– –
11.10 Write the resonance structures of CO3 and HCO3 .
2–
11.11 What is the state of hybridisation of carbon in (a) CO 3 (b) diamond
(c) graphite?
11.12 Explain the difference in properties of diamond and graphite on the basis
of their structures.
11.13 Rationalise the given statements and give chemical reactions :
• Lead(II) chloride reacts with Cl2 to give PbCl4.
• Lead(IV) chloride is highly unstable towards heat.
• Lead is known not to form an iodide, PbI4.
–
11.14 Suggest reasons why the B–F bond lengths in BF 3 (130 pm) and BF 4
(143 pm) differ.
11.15 If B–Cl bond has a dipole moment, explain why BCl3 molecule has zero
dipole moment.
11.16 Aluminium trifluoride is insoluble in anhydrous HF but dissolves on
addition of NaF. Aluminium trifluoride precipitates out of the resulting
solution when gaseous BF3 is bubbled through. Give reasons.
11.17 Suggest a reason as to why CO is poisonous.
11.18 How is excessive content of CO2 responsible for global warming ?
11.19 Explain structures of diborane and boric acid.
11.20 What happens when
(a) Borax is heated strongly,
(b) Boric acid is added to water,
(c) Aluminium is treated with dilute NaOH,
(d) BF3 is reacted with ammonia ?
11.21 Explain the following reactions
(a) Silicon is heated with methyl chloride at high temperature in the
presence of copper;
(b) Silicon dioxide is treated with hydrogen fluoride;
(c) CO is heated with ZnO;
(d) Hydrated alumina is treated with aqueous NaOH solution.
11.22 Give reasons :
(i) Conc. HNO3 can be transported in aluminium container.
(ii) A mixture of dilute NaOH and aluminium pieces is used to open
drain.
(iii) Graphite is used as lubricant.
(iv) Diamond is used as an abrasive.
(v) Aluminium alloys are used to make aircraft body.
(vi) Aluminium utensils should not be kept in water overnight.
(vii) Aluminium wire is used to make transmission cables.
11.23 Explain why is there a phenomenal decrease in ionization enthalpy from
carbon to silicon ?
11.24 How would you explain the lower atomic radius of Ga as compared to Al ?
11.25 What are allotropes? Sketch the structure of two allotropes of carbon namely
diamond and graphite. What is the impact of structure on physical
properties of two allotropes?
THE p-BLOCK ELEMENTS 325
UNIT 8
REDOX REACTIONS
In reactions (8.1) and (8.2), the elements been broadened these days to include removal
magnesium and sulphur are oxidised on of oxygen/electronegative element from a
account of addition of oxygen to them. substance or addition of hydrogen/
Similarly, methane is oxidised owing to the electropositive element to a substance.
addition of oxygen to it. According to the definition given above, the
CH4 (g) + 2O2 (g) → CO2 (g) + 2H2O (l) (8.3) following are the examples of reduction
processes:
A careful examination of reaction (8.3) in
which hydrogen has been replaced by oxygen 2 HgO (s) 2 Hg (l) + O2 (g) (8.8)
prompted chemists to reinterpret oxidation in (removal of oxygen from mercuric oxide )
terms of removal of hydrogen from it and,
2 FeCl3 (aq) + H2 (g) →2 FeCl2 (aq) + 2 HCl(aq)
therefore, the scope of term oxidation was
(8.9)
broadened to include the removal of hydrogen
from a substance. The following illustration is (removal of electronegative element, chlorine
another reaction where removal of hydrogen from ferric chloride)
can also be cited as an oxidation reaction. CH2 = CH2 (g) + H2 (g) → H3C – CH3 (g) (8.10)
2 H2S(g) + O2 (g) → 2 S (s) + 2 H2O (l) (8.4) (addition of hydrogen)
As knowledge of chemists grew, it was 2HgCl2 (aq) + SnCl2 (aq) → Hg2Cl2 (s)+SnCl4 (aq)
natural to extend the term oxidation for (8.11)
reactions similar to (8.1 to 8.4), which do not (addition of mercury to mercuric chloride)
involve oxygen but other electronegative In reaction (8.11) simultaneous oxidation
elements. The oxidation of magnesium with of stannous chloride to stannic chloride is also
fluorine, chlorine and sulphur etc. occurs occurring because of the addition of
according to the following reactions : electronegative element chlorine to it. It was
soon realised that oxidation and reduction
Mg (s) + F2 (g) → MgF2 (s) (8.5) always occur simultaneously (as will be
apparent by re-examining all the equations
Mg (s) + Cl2 (g) → MgCl2 (s) (8.6)
given above), hence, the word “redox” was
Mg (s) + S (s) → MgS (s) (8.7) coined for this class of chemical reactions.
Incorporating the reactions (8.5 to 8.7) Problem 8.1
within the fold of oxidation reactions In the reactions given below, identify the
encouraged chemists to consider not only the species undergoing oxidation and
removal of hydrogen as oxidation, but also the reduction:
removal of electropositive elements as
(i) H2S (g) + Cl2 (g) → 2 HCl (g) + S (s)
oxidation. Thus the reaction :
(ii) 3Fe3O4 (s) + 8 Al (s) → 9 Fe (s)
2K4 [Fe(CN)6](aq) + H2O2 (aq) →2K3[Fe(CN)6](aq)
+ 2 KOH (aq) + 4Al2O3 (s)
is interpreted as oxidation due to the removal (iii) 2 Na (s) + H2 (g) → 2 NaH (s)
of electropositive element potassium from Solution
potassium ferrocyanide before it changes to
(i) H 2 S is oxidised because a more
potassium ferricyanide. To summarise, the
electronegative element, chlorine is added
term “oxidation” is defined as the addition
to hydrogen (or a more electropositive
of oxygen/electronegative element to a
element, hydrogen has been removed
substance or removal of hydrogen/
from S). Chlorine is reduced due to
electropositive element from a substance.
addition of hydrogen to it.
In the beginning, reduction was
(ii) Aluminium is oxidised because
considered as removal of oxygen from a
oxygen is added to it. Ferrous ferric oxide
compound. However, the term reduction has
REDOX REACTIONS 257
(Fe3O4) is reduced because oxygen has For convenience, each of the above
been removed from it. processes can be considered as two separate
steps, one involving the loss of electrons and
(iii) With the careful application of the
the other the gain of electrons. As an
concept of electronegativity only we may
illustration, we may further elaborate one of
infer that sodium is oxidised and
these, say, the formation of sodium chloride.
hydrogen is reduced.
2 Na(s) → 2 Na (g) + 2e
+ –
Reaction (iii) chosen here prompts us to
think in terms of another way to define Cl2(g) + 2e → 2 Cl (g)
– –
redox reactions.
Each of the above steps is called a half
reaction, which explicitly shows involvement
8.2 REDOX REACTIONS IN TERMS OF
of electrons. Sum of the half reactions gives
ELECTRON TRANSFER REACTIONS
the overall reaction :
We have already learnt that the reactions
2 Na(s) + Cl2 (g) → 2 Na Cl (s) or 2 NaCl (s)
+ –
2Na(s) + Cl2(g) → 2NaCl (s) (8.12)
Reactions 8.12 to 8.14 suggest that half
2Na(s) + O2(g) → Na2O(s) (8.13)
reactions that involve loss of electrons are
2Na(s) + S(s) → Na2S(s) (8.14) called oxidation reactions. Similarly, the
are redox reactions because in each of these half reactions that involve gain of electrons
reactions sodium is oxidised due to the are called reduction reactions. It may not
addition of either oxygen or more be out of context to mention here that the new
electronegative element to sodium. way of defining oxidation and reduction has
Simultaneously, chlorine, oxygen and sulphur been achieved only by establishing a
are reduced because to each of these, the correlation between the behaviour of species
electropositive element sodium has been as per the classical idea and their interplay in
added. From our knowledge of chemical electron-transfer change. In reactions (8.12 to
bonding we also know that sodium chloride, 8.14) sodium, which is oxidised, acts as
sodium oxide and sodium sulphide are ionic a reducing agent because it donates electron
compounds and perhaps better written as to each of the elements interacting with it and
+ – + 2– + 2–
Na Cl (s), (Na ) 2O (s), and (Na ) 2 S (s). thus helps in reducing them. Chlorine, oxygen
Development of charges on the species and sulphur are reduced and act as oxidising
produced suggests us to rewrite the reactions agents because these accept electrons from
(8.12 to 8.14) in the following manner : sodium. To summarise, we may mention that
Oxidation : Loss of electron(s) by any species.
Reduction: Gain of electron(s) by any species.
Oxidising agent : Acceptor of electron(s).
Reducing agent : Donor of electron(s).
and the other half reaction is: At this stage we may investigate the state
H2 (g) + 2e → 2 H (g)
– – of equilibrium for the reaction represented by
equation (8.15). For this purpose, let us place
This splitting of the reaction under a strip of metallic copper in a zinc sulphate
examination into two half reactions solution. No visible reaction is noticed and
automatically reveals that here sodium is attempt to detect the presence of Cu2+ ions by
oxidised and hydrogen is reduced, passing H 2S gas through the solution to
therefore, the complete reaction is a redox produce the black colour of cupric sulphide,
change. CuS, does not succeed. Cupric sulphide has
such a low solubility that this is an extremely
8.2.1 Competitive Electron Transfer
sensitive test; yet the amount of Cu2+ formed
Reactions
cannot be detected. We thus conclude that the
Place a strip of metallic zinc in an aqueous
state of equilibrium for the reaction (8.15)
solution of copper nitrate as shown in Fig. 8.1,
greatly favours the products over the reactants.
for about one hour. You may notice that the
strip becomes coated with reddish metallic Let us extend electron transfer reaction now
copper and the blue colour of the solution to copper metal and silver nitrate solution in
2+
disappears. Formation of Zn ions among the water and arrange a set-up as shown in
products can easily be judged when the blue Fig. 8.2. The solution develops blue colour due
colour of the solution due to Cu has
2+ to the formation of Cu2+ ions on account of the
disappeared. If hydrogen sulphide gas is reaction:
passed through the colourless solution
2+
containing Zn ions, appearance of white zinc
sulphide, ZnS can be seen on making the
solution alkaline with ammonia.
The reaction between metallic zinc and the (8.16)
aqueous solution of copper nitrate is : 2+
Here, Cu(s) is oxidised to Cu (aq) and
Zn(s) + Cu (aq) → Zn (aq) + Cu(s)
2+ 2+ +
(8.15) Ag (aq) is reduced to Ag(s). Equilibrium greatly
2+
In reaction (8.15), zinc has lost electrons favours the products Cu (aq) and Ag(s).
2+
to form Zn and, therefore, zinc is oxidised. By way of contrast, let us also compare the
Evidently, now if zinc is oxidised, releasing reaction of metallic cobalt placed in nickel
electrons, something must be reduced, sulphate solution. The reaction that occurs
accepting the electrons lost by zinc. Copper here is :
ion is reduced by gaining electrons from the zinc.
Reaction (8.15) may be rewritten as :
(8.17)
Fig. 8.1 Redox reaction between zinc and aqueous solution of copper nitrate occurring in a beaker.
REDOX REACTIONS 259
Fig. 8.2 Redox reaction between copper and aqueous solution of silver nitrate occurring in a beaker.
At equilibrium, chemical tests reveal that both However, as we shall see later, the charge
2+ 2+
Ni (aq) and Co (aq) are present at moderate transfer is only partial and is perhaps better
concentrations. In this case, neither the described as an electron shift rather than a
reactants [Co(s) and Ni2+(aq)] nor the products complete loss of electron by H and gain by O.
2+
[Co (aq) and Ni (s)] are greatly favoured. What has been said here with respect to
This competition for release of electrons equation (8.18) may be true for a good number
incidently reminds us of the competition for of other reactions involving covalent
release of protons among acids. The similarity compounds. Two such examples of this class
suggests that we might develop a table in of the reactions are:
which metals and their ions are listed on the H2(s) + Cl2(g) → 2HCl(g) (8.19)
basis of their tendency to release electrons just and,
as we do in the case of acids to indicate the
CH 4(g) + 4Cl2(g) → CCl4(l) + 4HCl(g) (8.20)
strength of the acids. As a matter of fact we
have already made certain comparisons. By In order to keep track of electron shifts in
comparison we have come to know that zinc chemical reactions involving formation of
releases electrons to copper and copper covalent compounds, a more practical method
releases electrons to silver and, therefore, the of using oxidation number has been
electron releasing tendency of the metals is in developed. In this method, it is always
the order: Zn>Cu>Ag. We would love to make assumed that there is a complete transfer of
our list more vast and design a metal activity electron from a less electronegative atom to a
series or electrochemical series. The more electonegative atom. For example, we
competition for electrons between various rewrite equations (8.18 to 8.20) to show
charge on each of the atoms forming part of
metals helps us to design a class of cells,
the reaction :
named as Galvanic cells in which the chemical
0 0 +1 –2
reactions become the source of electrical
energy. We would study more about these cells 2H2(g) + O2(g) → 2H2O (l) (8.21)
in Class XII. 0 0 +1 –1
H2 (s) + Cl2(g) → 2HCl(g) (8.22)
8.3 OXIDATION NUMBER
– 4 +1 0 +4 –1 +1 –1
A less obvious example of electron transfer is
CH4(g) + 4Cl2(g) → 4CCl4(l) +4HCl(g) (8.23)
realised when hydrogen combines with oxygen
to form water by the reaction: It may be emphasised that the assumption
2H2(g) + O2 (g) → 2H2O (l) (8.18) of electron transfer is made for book-keeping
Though not simple in its approach, yet we purpose only and it will become obvious at a
can visualise the H atom as going from a later stage in this unit that it leads to the simple
neutral (zero) state in H2 to a positive state in description of redox reactions.
H2O, the O atom goes from a zero state in O2 Oxidation number denotes the
to a dinegative state in H2O. It is assumed that oxidation state of an element in a
there is an electron transfer from H to O and compound ascertained according to a set
consequently H2 is oxidised and O2 is reduced. of rules formulated on the basis that
260 CHEMISTRY
electron in a covalent bond belongs of oxygen but this number would now be
entirely to more electronegative element. a positive figure only.
It is not always possible to remember or 4. The oxidation number of hydrogen is +1,
make out easily in a compound/ion, which except when it is bonded to metals in binary
element is more electronegative than the other. compounds (that is compounds containing
Therefore, a set of rules has been formulated two elements). For example, in LiH, NaH,
to determine the oxidation number of an and CaH2, its oxidation number is –1.
element in a compound/ion. If two or more 5. In all its compounds, fluorine has an
than two atoms of an element are present in oxidation number of –1. Other halogens (Cl,
2–
the molecule/ion such as Na2S2O3/Cr2O7 , the Br, and I) also have an oxidation number
oxidation number of the atom of that element of –1, when they occur as halide ions in
will then be the average of the oxidation their compounds. Chlorine, bromine and
number of all the atoms of that element. We iodine when combined with oxygen, for
may at this stage, state the rules for the example in oxoacids and oxoanions, have
calculation of oxidation number. These rules are: positive oxidation numbers.
6. The algebraic sum of the oxidation number
1. In elements, in the free or the uncombined
of all the atoms in a compound must be
state, each atom bears an oxidation
zero. In polyatomic ion, the algebraic sum
number of zero. Evidently each atom in H2,
of all the oxidation numbers of atoms of
O2, Cl2, O3, P4, S8, Na, Mg, Al has the
the ion must equal the charge on the ion.
oxidation number zero.
Thus, the sum of oxidation number of three
2. For ions composed of only one atom, the oxygen atoms and one carbon atom in the
oxidation number is equal to the charge carbonate ion, (CO3)2– must equal –2.
+
on the ion. Thus Na ion has an oxidation By the application of above rules, we can
2+ 3+
number of +1, Mg ion, +2, Fe ion, +3, find out the oxidation number of the desired
– 2–
Cl ion, –1, O ion, –2; and so on. In their element in a molecule or in an ion. It is clear
compounds all alkali metals have that the metallic elements have positive
oxidation number of +1, and all alkaline oxidation number and nonmetallic elements
earth metals have an oxidation number of have positive or negative oxidation number.
+2. Aluminium is regarded to have an The atoms of transition elements usually
oxidation number of +3 in all its display several positive oxidation states. The
compounds. highest oxidation number of a representative
3. The oxidation number of oxygen in most element is the group number for the first two
compounds is –2. However, we come across groups and the group number minus 10
two kinds of exceptions here. One arises (following the long form of periodic table) for
in the case of peroxides and superoxides, the other groups. Thus, it implies that the
the compounds of oxygen in which oxygen highest value of oxidation number exhibited
atoms are directly linked to each other. by an atom of an element generally increases
While in peroxides (e.g., H2O2, Na2O2), each across the period in the periodic table. In the
oxygen atom is assigned an oxidation third period, the highest value of oxidation
number of –1, in superoxides (e.g., KO2, number changes from 1 to 7 as indicated below
RbO2) each oxygen atom is assigned an in the compounds of the elements.
oxidation number of –(½). The second A term that is often used interchangeably
exception appears rarely, i.e. when oxygen with the oxidation number is the oxidation
is bonded to fluorine. In such compounds state. Thus in CO2, the oxidation state of
e.g., oxygen difluoride (OF2) and dioxygen carbon is +4, that is also its oxidation number
difluoride (O2F2), the oxygen is assigned and similarly the oxidation state as well as
an oxidation number of +2 and +1, oxidation number of oxygen is – 2. This implies
respectively. The number assigned to that the oxidation number denotes the
oxygen will depend upon the bonding state oxidation state of an element in a compound.
REDOX REACTIONS 261
Group 1 2 13 14 15 16 17
Element Na Mg Al Si P S Cl
Compound NaCl MgSO4 AlF3 SiCl4 P4O10 SF6 HClO4
Highest oxidation +1 +2 +3 +4 +5 +6 +7
number state of
the group element
The oxidation number state of a metal in a The idea of oxidation number has been
compound is sometimes presented according invariably applied to define oxidation,
to the notation given by German chemist, reduction, oxidising agent (oxidant), reducing
Alfred Stock. It is popularly known as Stock agent (reductant) and the redox reaction. To
notation. According to this, the oxidation summarise, we may say that:
number is expressed by putting a Roman
Oxidation: An increase in the oxidation
numeral representing the oxidation number
number of the element in the given substance.
in parenthesis after the symbol of the metal in
the molecular formula. Thus aurous chloride Reduction: A decrease in the oxidation
and auric chloride are written as Au(I)Cl and number of the element in the given substance.
Au(III)Cl3. Similarly, stannous chloride and Oxidising agent: A reagent which can
stannic chloride are written as Sn(II)Cl2 and increase the oxidation number of an element
Sn(IV)Cl4. This change in oxidation number in a given substance. These reagents are called
implies change in oxidation state, which in as oxidants also.
turn helps to identify whether the species is Reducing agent: A reagent which lowers the
present in oxidised form or reduced form. oxidation number of an element in a given
Thus, Hg2(I)Cl2 is the reduced form of Hg(II) Cl2. substance. These reagents are also called as
Problem 8.3 reductants.
Using Stock notation, represent the Redox reactions: Reactions which involve
following compounds :HAuCl4, Tl2O, FeO, change in oxidation number of the interacting
Fe2O3, CuI, CuO, MnO and MnO2. species.
Solution Problem 8.4
By applying various rules of calculating Justify that the reaction:
the oxidation number of the desired
element in a compound, the oxidation 2Cu2O(s) + Cu2S(s) → 6Cu(s) + SO2(g)
number of each metallic element in its is a redox reaction. Identify the species
compound is as follows: oxidised/reduced, which acts as an
HAuCl4 → Au has 3 oxidant and which acts as a reductant.
Tl2O → Tl has 1 Solution
FeO → Fe has 2 Let us assign oxidation number to each
Fe2O3 → Fe has 3 of the species in the reaction under
CuI → Cu has 1 examination. This results into:
CuO → Cu has 2 +1 –2 +1 –2 0 +4 –2
MnO → Mn has 2 2Cu2O(s) + Cu2S(s) → 6Cu(s) + SO2
MnO2 → Mn has 4
We therefore, conclude that in this
Therefore, these compounds may be
reaction copper is reduced from +1 state
represented as:
to zero oxidation state and sulphur is
HAu(III)Cl 4, Tl2(I)O, Fe(II)O, Fe 2(III)O3, oxidised from –2 state to +4 state. The
Cu(I)I, Cu(II)O, Mn(II)O, Mn(IV)O2. above reaction is thus a redox reaction.
262 CHEMISTRY
Further, Cu2O helps sulphur in Cu2S to that all decomposition reactions are not redox
increase its oxidation number, therefore, reactions. For example, decomposition of
Cu(I) is an oxidant; and sulphur of Cu2S calcium carbonate is not a redox reaction.
helps copper both in Cu2S itself and Cu2O +2 +4 –2 +2 –2 +4 –2
to decrease its oxidation number; CaCO3 (s) CaO(s) + CO2(g)
therefore, sulphur of Cu2S is reductant.
3. Displacement reactions
8.3.1 Types of Redox Reactions In a displacement reaction, an ion (or an atom)
in a compound is replaced by an ion (or an
1. Combination reactions atom) of another element. It may be denoted
A combination reaction may be denoted in the as:
manner: X + YZ → XZ + Y
A+B → C
Displacement reactions fit into two categories:
Either A and B or both A and B must be in the metal displacement and non-metal
elemental form for such a reaction to be a redox displacement.
reaction. All combustion reactions, which
make use of elemental dioxygen, as well as (a) Metal displacement: A metal in a
other reactions involving elements other than compound can be displaced by another metal
dioxygen, are redox reactions. Some important in the uncombined state. We have already
examples of this category are: discussed about this class of the reactions
under section 8.2.1. Metal displacement
0 0 +4 –2
C(s) + O2 (g) CO2(g) (8.24) reactions find many applications in
metallurgical processes in which pure metals
0 0 +2 –3 are obtained from their compounds in ores. A
3Mg(s) + N2(g) Mg3N2(s) (8.25) few such examples are:
–4+1 0 +4 –2 +1 –2 +2 +4 –2 0 0 +2 +4 –2
CuSO4(aq) + Zn (s) → Cu(s) + ZnSO4 (aq)
CH4(g) + 2O2(g) CO2(g) + 2H2O (l)
(8.29)
2. Decomposition reactions +5 –2 0 0 +2 –2
Decomposition reactions are the opposite of V2O5 (s) + 5Ca (s) 2V (s) + 5CaO (s)
combination reactions. Precisely, a (8.30)
decomposition reaction leads to the breakdown
of a compound into two or more components +4 –1 0 0 +2 –1
at least one of which must be in the elemental TiCl4 (l) + 2Mg (s) Ti (s) + 2 MgCl2 (s)
state. Examples of this class of reactions are: (8.31)
+1 –2 0 0 +3 –2 0 +3 –2 0
2H2O (l) 2H2 (g) + O2(g) (8.26) Cr2O3 (s) + 2 Al (s) Al2O3 (s) + 2Cr(s)
+1 –1 0 0 (8.32)
2NaH (s) 2Na (s) + H2(g) (8.27) In each case, the reducing metal is a better
reducing agent than the one that is being
+1 +5 –2 +1 –1 0 reduced which evidently shows more capability
2KClO3 (s) 2KCl (s) + 3O2(g) (8.28) to lose electrons as compared to the one that
It may carefully be noted that there is no is reduced.
change in the oxidation number of hydrogen (b) Non-metal displacement: The non-metal
in methane under combination reactions and displacement redox reactions include
that of potassium in potassium chlorate in hydrogen displacement and a rarely occurring
reaction (8.28). This may also be noted here reaction involving oxygen displacement.
REDOX REACTIONS 263
All alkali metals and some alkaline earth order Zn> Cu>Ag. Like metals, activity series
metals (Ca, Sr, and Ba) which are very good also exists for the halogens. The power of these
reductants, will displace hydrogen from cold elements as oxidising agents decreases as we
water. move down from fluorine to iodine in group
0 +1 –2 +1 –2 +1 0 17 of the periodic table. This implies that
2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g) fluorine is so reactive that it can replace
(8.33) chloride, bromide and iodide ions in solution.
0 +1 –2 +2 –2 +1 0 In fact, fluorine is so reactive that it attacks
Ca(s) + 2H2O(l) → Ca(OH)2 (aq) + H2(g) water and displaces the oxygen of water :
(8.34) +1 –2 0 +1 –1 0
Less active metals such as magnesium and 2H2O (l) + 2F2 (g) → 4HF(aq) + O2(g) (8.40)
iron react with steam to produce dihydrogen gas: It is for this reason that the displacement
0 +1 –2 +2 –2 +1 0 reactions of chlorine, bromine and iodine
Mg(s) + 2H2O(l) Mg(OH)2(s) + H2(g) using fluorine are not generally carried out in
(8.35) aqueous solution. On the other hand, chlorine
can displace bromide and iodide ions in an
0 +1 –2 +3 –2 0
aqueous solution as shown below:
2Fe(s) + 3H2O(l) Fe2O3(s) + 3H2(g) (8.36)
0 +1 –1 +1 –1 0
Many metals, including those which do not Cl2 (g) + 2KBr (aq) → 2 KCl (aq) + Br2 (l)
react with cold water, are capable of displacing (8.41)
hydrogen from acids. Dihydrogen from acids 0 +1–1 +1 –1 0
may even be produced by such metals which Cl2 (g) + 2KI (aq) → 2 KCl (aq) + I2 (s)
do not react with steam. Cadmium and tin are (8.42)
the examples of such metals. A few examples As Br2 and I2 are coloured and dissolve in CCl4,
for the displacement of hydrogen from acids can easily be identified from the colour of the
are: solution. The above reactions can be written
0 +1 –1 +2 –1 0 in ionic form as:
Zn(s) + 2HCl(aq) → ZnCl2 (aq) + H2 (g) 0 –1 –1 0
Cl2 (g) + 2Br (aq) → 2Cl (aq) + Br2 (l) (8.41a)
– –
(8.37)
0 +1 –1 +2 –1 0 0 –1 –1 0
Mg (s) + 2HCl (aq) → MgCl2 (aq) + H2 (g) Cl2 (g) + 2I (aq) → 2Cl (aq) + I2 (s) (8.42b)
– –
(8.38)
Reactions (8.41) and (8.42) form the basis
0 +1 –1 +2 –1 0 – –
of identifying Br and I in the laboratory
Fe(s) + 2HCl(aq) → FeCl2(aq) + H2(g)
through the test popularly known as ‘Layer
(8.39)
Test’. It may not be out of place to mention
Reactions (8.37 to 8.39) are used to here that bromine likewise can displace iodide
prepare dihydrogen gas in the laboratory. ion in solution:
Here, the reactivity of metals is reflected in the
0 –1 –1 0
rate of hydrogen gas evolution, which is the
Br2 (l) + 2I (aq) → 2Br (aq) + I2 (s)
– –
(8.43)
slowest for the least active metal Fe, and the
fastest for the most reactive metal, Mg. Very The halogen displacement reactions have
less active metals, which may occur in the a direct industrial application. The recovery
native state such as silver (Ag), and gold (Au) of halogens from their halides requires an
do not react even with hydrochloric acid. oxidation process, which is represented by:
→ X2 + 2e
– –
In section (8.2.1) we have already 2X (8.44)
discussed that the metals – zinc (Zn), copper here X denotes a halogen element. Whereas
–
(Cu) and silver (Ag) through tendency to lose chemical means are available to oxidise Cl ,
– –
electrons show their reducing activity in the Br and I , as fluorine is the strongest oxidising
264 CHEMISTRY
–
agent; there is no way to convert F ions to F2 fluorine shows deviation from this behaviour
by chemical means. The only way to achieve when it reacts with alkali. The reaction that
F2 from F– is to oxidise electrolytically, the takes place in the case of fluorine is as follows:
details of which you will study at a later stage. 2 F2(g) + 2OH (aq) → 2 F (aq) + OF2(g) + H2O(l)
– –
+ 6H2O(l)
+5 –1 +7
(8.47) – – –
4ClO 3 → Cl + 3 ClO4
0 +1 –1
Cl2 (g) + 2 OH– (aq) → ClO– (aq) + Cl– (aq) + Problem 8.6
H2O (l) Suggest a scheme of classification of the
(8.48) following redox reactions
The reaction (8.48) describes the formation
(a) N2 (g) + O2 (g) → 2 NO (g)
of household bleaching agents. The
hypochlorite ion (ClO – ) formed in the reaction (b) 2Pb(NO3)2(s) → 2PbO(s) + 2 NO2 (g) +
oxidises the colour -bearing stains of the ½ O2 (g)
substances to colourless compounds. (c) NaH(s) + H2O(l) → NaOH(aq) + H2 (g)
– –
It is of interest to mention here that whereas (d) 2NO2(g) + 2OH (aq) → NO2(aq) +
–
bromine and iodine follow the same trend as NO3 (aq)+H2O(l)
exhibited by chlorine in reaction (8.48),
REDOX REACTIONS 265
+6 –2 +4 –2 +3 +6 –2
2– 2–
Cr2O7 (aq) + SO3 (aq) → Cr(aq)+SO4 (aq)
2– Step 4: As the reaction occurs in the basic
medium, and the ionic charges are not
This indicates that the dichromate ion is equal on both sides, add 2 OH ions on
–
the oxidant and the sulphite ion is the the right to make ionic charges equal.
reductant. – –
2MnO4 (aq) + Br (aq) → 2MnO2(s) +
Step 3: Calculate the increase and – –
decrease of oxidation number, and make BrO3 (aq) + 2OH (aq)
them equal: Step 5: Finally, count the hydrogen
+6 –2 +4 –2 +3
atoms and add appropriate number of
Cr2O7 (aq) + 3SO3 (aq) → 2Cr
2– 2– 3+
(aq) + water molecules (i.e. one H2O molecule)
+6 –2
on the left side to achieve balanced redox
3SO4 (aq)
2– change.
– –
Step 4: As the reaction occurs in the 2MnO4(aq) + Br (aq) + H2O(l) → 2MnO2(s)
– –
acidic medium, and further the ionic + BrO3 (aq) + 2OH (aq)
charges are not equal on both the sides, (b) Half Reaction Method: In this method,
+
add 8H on the left to make ionic charges the two half equations are balanced separately
equal and then added together to give balanced
2– 2– + 3+
Cr2O7 (aq) + 3SO3 (aq)+ 8H → 2Cr (aq) equation.
2–
+ 3SO4 (aq) Suppose we are to balance the equation
2+ 3+
Step 5: Finally, count the hydrogen showing the oxidation of Fe ions to Fe ions
2–
atoms, and add appropriate number of by dichromate ions (Cr2O7) in acidic medium,
2– 3+
water molecules (i.e., 4H2O) on the right wherein, Cr2O7 ions are reduced to Cr ions.
to achieve balanced redox change. The following steps are involved in this task.
2– 2– +
Cr2O7 (aq) + 3SO3 (aq)+ 8H (aq) → Step 1: Produce unbalanced equation for the
3+ 2–
2Cr (aq) + 3SO4 (aq) +4H2O (l) reaction in ionic form :
2+ 2– 3+ 3+
Problem 8.9 Fe (aq) + Cr2O7 (aq) → Fe (aq) + Cr (aq)
(8.50)
Permanganate ion reacts with bromide ion Step 2: Separate the equation into half-
in basic medium to give manganese reactions:
dioxide and bromate ion. Write the
+2 +3
balanced ionic equation for the reaction.
(aq) → Fe (aq)
2+ 3+
Oxidation half : Fe (8.51)
Solution +6 –2 +3
2– 3+
Step 1 : The skeletal ionic equation is : Reduction half : Cr2O7 (aq) → Cr (aq)
– – –
MnO4 (aq) + Br (aq) → MnO2(s) + BrO3 (aq) (8.52)
Step 2 : Assign oxidation numbers for Step 3: Balance the atoms other than O and
Mn and Br H in each half reaction individually. Here the
oxidation half reaction is already balanced with
+7 –1 +4 +5
– – – respect to Fe atoms. For the reduction half
MnO4(aq) + Br (aq) →MnO2 (s) + BrO3 (aq) 3+
reaction, we multiply the Cr by 2 to balance
this indicates that permanganate ion is Cr atoms.
268 CHEMISTRY
2– 3+
Cr2O7 (aq) → 2 Cr (aq) (8.53) Problem 8.10
Step 4: For reactions occurring in acidic –
Permanganate(VII) ion, MnO4 in basic
+
medium, add H2O to balance O atoms and H –
solution oxidises iodide ion, I to produce
to balance H atoms. molecular iodine (I2) and manganese (IV)
Thus, we get : oxide (MnO2). Write a balanced ionic
2– + 3+ equation to represent this redox reaction.
Cr2O7 (aq) + 14H (aq) → 2 Cr (aq) + 7H2O (l)
(8.54) Solution
Step 5: Add electrons to one side of the half Step 1: First we write the skeletal ionic
reaction to balance the charges. If need be, equation, which is
–
MnO4 (aq) + I (aq) → MnO2(s) + I2(s)
–
make the number of electrons equal in the two
half reactions by multiplying one or both half Step 2: The two half-reactions are:
reactions by appropriate coefficients.
–1 0
The oxidation half reaction is thus rewritten Oxidation half : I (aq) → I2 (s)
–
To equalise the number of electrons in both To balance the H atoms, we add four H
+
– –
6I (aq) → 3I2 (s) + 6e (iii) There is yet another method which is
– – interesting and quite common. Its use is
2 MnO4 (aq) + 4H2O (l) +6e → 2MnO2(s)
– restricted to those reagents which are able
+ 8OH (aq) –
to oxidise I ions, say, for example, Cu(II):
Step 6: Add two half-reactions to obtain –
2Cu2+(aq) + 4I (aq) → Cu2I2(s) + I2(aq) (8.59)
the net reactions after cancelling electrons
on both sides. This method relies on the facts that iodine
– – itself gives an intense blue colour with starch
6I (aq) + 2MnO4(aq) + 4H2O(l) → 3I2(s) +
– and has a very specific reaction with
2MnO2(s) +8 OH (aq)
thiosulphate ions (S2O32–), which too is a redox
Step 7: A final verification shows that reaction:
the equation is balanced in respect of the 2–
I2(aq) + 2 S2O3 (aq)→2I–(aq) + S4O62–(aq) (8.60)
number of atoms and charges on both
sides. I2, though insoluble in water, remains in
solution containing KI as KI3.
8.3.3 Redox Reactions as the Basis for On addition of starch after the liberation of
Titrations iodine from the reaction of Cu2+ ions on iodide
In acid-base systems we come across with a ions, an intense blue colour appears. This
titration method for finding out the strength colour disappears as soon as the iodine is
of one solution against the other using a pH consumed by the thiosulphate ions. Thus, the
sensitive indicator. Similarly, in redox systems, end-point can easily be tracked and the rest
the titration method can be adopted to is the stoichiometric calculation only.
determine the strength of a reductant/oxidant 8.3.4 Limitations of Concept of Oxidation
using a redox sensitive indicator. The usage Number
of indicators in redox titration is illustrated
As you have observed in the above discussion,
below:
the concept of redox processes has been
(i) In one situation, the reagent itself is evolving with time. This process of evolution
intensely coloured, e.g., permanganate ion, is continuing. In fact, in recent past the
– –
MnO4. Here MnO4 acts as the self indicator. oxidation process is visualised as a decrease
The visible end point in this case is in electron density and reduction process as
2+
achieved after the last of the reductant (Fe an increase in electron density around the
2–
or C2O4 ) is oxidised and the first lasting atom(s) involved in the reaction.
–
tinge of pink colour appears at MnO4
–6 –3
concentration as low as 10 mol dm 8.4 REDOX REACTIONS AND ELECTRODE
–6 –1
(10 mol L ). This ensures a minimal PROCESSES
‘overshoot’ in colour beyond the The experiment corresponding to reaction
equivalence point, the point where the (8.15), can also be observed if zinc rod is
reductant and the oxidant are equal in dipped in copper sulphate solution. The redox
terms of their mole stoichiometry. reaction takes place and during the reaction,
(ii) If there is no dramatic auto-colour change zinc is oxidised to zinc ions and copper ions
–
(as with MnO 4 titration), there are are reduced to metallic copper due to direct
indicators which are oxidised immediately transfer of electrons from zinc to copper ion.
after the last bit of the reactant is During this reaction heat is also evolved. Now
consumed, producing a dramatic colour we modify the experiment in such a manner
change. The best example is afforded by that for the same redox reaction transfer of
–
Cr2O72 , which is not a self-indicator, but electrons takes place indirectly. This
oxidises the indicator substance necessitates the separation of zinc metal from
diphenylamine just after the equivalence copper sulphate solution. We take copper
point to produce an intense blue colour, sulphate solution in a beaker and put a copper
thus signalling the end point. strip or rod in it. We also take zinc sulphate
270 CHEMISTRY
solution in another beaker and put a zinc rod jelly like substance). This provides an electric
or strip in it. Now reaction takes place in either contact between the two solutions without
of the beakers and at the interface of the metal allowing them to mix with each other. The
and its salt solution in each beaker both the zinc and copper rods are connected by a metallic
reduced and oxidized forms of the same wire with a provision for an ammeter and a
species are present. These represent the switch. The set-up as shown in Fig.8.3 is known
species in the reduction and oxidation half as Daniell cell. When the switch is in the off
reactions. A redox couple is defined as having position, no reaction takes place in either of
together the oxidised and reduced forms of a the beakers and no current flows through the
substance taking part in an oxidation or metallic wire. As soon as the switch is in the
reduction half reaction. on position, we make the following
This is represented by separating the observations:
oxidised form from the reduced form by a 1. The transfer of electrons now does not take
2+
vertical line or a slash representing an place directly from Zn to Cu but through
interface (e.g. solid/solution). For example the metallic wire connecting the two rods
in this experiment the two redox couples are as is apparent from the arrow which
represented as Zn2+/Zn and Cu2+/Cu. In both indicates the flow of current.
cases, oxidised form is put before the reduced 2. The electricity from solution in one beaker
form. Now we put the beaker containing to solution in the other beaker flows by the
copper sulphate solution and the beaker migration of ions through the salt bridge.
containing zinc sulphate solution side by side We know that the flow of current is possible
(Fig. 8.3). We connect solutions in two only if there is a potential difference
beakers by a salt bridge (a U-tube containing between the copper and zinc rods known
a solution of potassium chloride or as electrodes here.
ammonium nitrate usually solidified by The potential associated with each
boiling with agar agar and later cooling to a electrode is known as electrode potential. If
the concentration of each species taking part
in the electrode reaction is unity (if any gas
appears in the electrode reaction, it is confined
to 1 atmospheric pressure) and further the
reaction is carried out at 298K, then the
potential of each electrode is said to be the
Standard Electrode Potential. By
convention, the standard electrode potential
0
(E of hydrogen electrode is 0.00 volts. The
electrode potential value for each electrode
process is a measure of the relative tendency
of the active species in the process to remain
0
in the oxidised/reduced form. A negative E
means that the redox couple is a stronger
+
reducing agent than the H /H2 couple. A
0
positive E means that the redox couple is a
Fig.8.3 The set-up for Daniell cell. Electrons +
weaker reducing agent than the H /H2 couple.
produced at the anode due to oxidation The standard electrode potentials are very
of Zn travel through the external circuit
to the cathode where these reduce the
important and we can get a lot of other useful
copper ions. The circuit is completed information from them. The values of standard
inside the cell by the migration of ions electrode potentials for some selected electrode
through the salt bridge. It may be noted processes (reduction reactions) are given in
that the direction of current is opposite Table 8.1. You will learn more about electrode
to the direction of electron flow. reactions and cells in Class XII.
REDOX REACTIONS 271
SUMMARY
Redox reactions form an important class of reactions in which oxidation and reduction
occur simultaneously. Three tier conceptualisation viz, classical, electronic and oxidation
number, which is usually available in the texts, has been presented in detail. Oxidation,
reduction, oxidising agent (oxidant) and reducing agent (reductant) have been viewed
according to each conceptualisation. Oxidation numbers are assigned in accordance
with a consistent set of rules. Oxidation number and ion-electron method both are
useful means in writing equations for the redox reactions. Redox reactions are classified
into four categories: combination, decomposition displacement and disproportionation
reactions. The concept of redox couple and electrode processes is introduced here.
The redox reactions find wide applications in the study of electrode processes and cells.
EXERCISES
8.1 Assign oxidation number to the underlined elements in each of the following
species:
(a) NaH2PO4 (b) NaHSO4 (c) H4P2O7 (d) K2MnO4
(e) CaO2 (f) NaBH4 (g) H2S2O7 (h) KAl(SO4)2.12 H2O
8.2 What are the oxidation number of the underlined elements in each of the
following and how do you rationalise your results ?
(a) KI3 (b) H2S4O6 (c) Fe3O4 (d) CH3CH2OH (e) CH3COOH
8.3 Justify that the following reactions are redox reactions:
(a) CuO(s) + H2(g) → Cu(s) + H2O(g)
(b) Fe2O3(s) + 3CO(g) → 2Fe(s) + 3CO2(g)
(c) 4BCl3(g) + 3LiAlH4(s) → 2B2H6(g) + 3LiCl(s) + 3 AlCl3 (s)
(d) 2K(s) + F2(g) → 2K F (s)
+ –
+ 2H2O(l)
(c) HCHO (l) + 2 Cu (aq) + 5 OH (aq) → Cu2O(s) + HCOO (aq) + 3H2O(l)
2+ – –
S2O3 (aq) + 2Br2(l) + 5 H2O(l) → 2SO4 (aq) + 4Br (aq) + 10H (aq)
2– 2– – +
Why does the same reductant, thiosulphate react differently with iodine and
bromine ?
8.15 Justify giving reactions that among halogens, fluorine is the best oxidant and
among hydrohalic compounds, hydroiodic acid is the best reductant.
8.16 Why does the following reaction occur ?
XeO6 (aq) + 2F (aq) + 6H (aq) → XeO3(g)+ F2(g) + 3H2O(l)
4– – +
4–
What conclusion about the compound Na4XeO6 (of which XeO6 is a part) can be
drawn from the reaction.
8.17 Consider the reactions:
(a) H3PO2(aq) + 4 AgNO3(aq) + 2 H2O(l) → H3PO4(aq) + 4Ag(s) + 4HNO3(aq)
(b) H3PO2(aq) + 2CuSO4(aq) + 2 H2O(l) → H3PO4(aq) + 2Cu(s) + H2SO4(aq)
(c) C6H5CHO(l) + 2[Ag (NH3)2] (aq) + 3OH (aq) → C6H5COO (aq) + 2Ag(s) +
+ – –
+ 2+
What inference do you draw about the behaviour of Ag and Cu from these
reactions ?
8.18 Balance the following redox reactions by ion – electron method :
(a) MnO4 (aq) + I (aq) → MnO2 (s) + I2(s) (in basic medium)
– –
8.20 What sorts of informations can you draw from the following reaction ?
(CN)2(g) + 2OH (aq) → CN (aq) + CNO (aq) + H2O(l)
– – –
3+
8.21 The Mn ion is unstable in solution and undergoes disproportionation to give
2+ +
Mn , MnO2, and H ion. Write a balanced ionic equation for the reaction.
8.22 Consider the elements :
Cs, Ne, I and F
(a) Identify the element that exhibits only negative oxidation state.
(b) Identify the element that exhibits only postive oxidation state.
(c) Identify the element that exhibits both positive and negative oxidation states.
(d) Identify the element which exhibits neither the negative nor does the positive
oxidation state.
8.23 Chlorine is used to purify drinking water. Excess of chlorine is harmful. The
excess of chlorine is removed by treating with sulphur dioxide. Present a balanced
equation for this redox change taking place in water.
8.24 Refer to the periodic table given in your book and now answer the following
questions:
(a) Select the possible non metals that can show disproportionation reaction.
(b) Select three metals that can show disproportionation reaction.
8.25 In Ostwald’s process for the manufacture of nitric acid, the first step involves
the oxidation of ammonia gas by oxygen gas to give nitric oxide gas and steam.
What is the maximum weight of nitric oxide that can be obtained starting only
with 10.00 g. of ammonia and 20.00 g of oxygen ?
8.26 Using the standard electrode potentials given in the Table 8.1, predict if the
reaction between the following is feasible:
3+ –
(a) Fe (aq) and I (aq)
+
(b) Ag (aq) and Cu(s)
3+
(c) Fe (aq) and Cu(s)
3+
(d) Ag(s) and Fe (aq)
2+
(e) Br2(aq) and Fe (aq).
REDOX REACTIONS 275
UNIT 12
(b)
Ethane Ethene
(c)
Ethyne Methanol
Solution
Condensed formula:
Cyclopropane (a) HO(CH2)3CH(CH3)CH(CH3)2
(b) HOCH(CN)2
Bond-line formula:
(a)
Cyclopentane
(b)
chlorocyclohexane
Problem 12.6
Problem 12.4
Expand each of the following bond-line
Expand each of the following condensed formulas to show all the atoms including
formulas into their complete structural carbon and hydrogen
formulas. (a)
(a) CH3CH2COCH2CH3
(b) CH3CH=CH(CH2)3CH3
Solution (b)
(a)
(c)
(b) (d)
Solution
Problem 12.5
For each of the following compounds,
write a condensed formula and also their
bond-line formula.
(a) HOCH2CH2CH2CH(CH3)CH(CH3)CH3
330 CHEMISTRY
Molecular Models
Molecular models are physical devices that
are used for a better visualisation and
perception of three-dimensional shapes of
organic molecules. These are made of wood,
plastic or metal and are commercially
available. Commonly three types of molecular
models are used: (1) Framework model, (2)
Ball-and-stick model, and (3) Space filling
model. In the framework model only the
bonds connecting the atoms of a molecule
and not the atoms themselves are shown.
This model emphasizes the pattern of bonds
of a molecule while ignoring the size of atoms.
In the ball-and-stick model, both the atoms
and the bonds are shown. Balls represent
atoms and the stick denotes a bond.
Compounds containing C=C (e.g., ethene) can
best be represented by using springs in place
12.3.2 Three-Dimensional
of sticks. These models are referred to as ball-
Representation of Organic and-spring model. The space-filling model
Molecules emphasises the relative size of each atom
The three-dimensional (3-D) structure of based on its van der Waals radius. Bonds
organic molecules can be represented on are not shown in this model. It conveys the
paper by using certain conventions. For volume occupied by each atom in the
example, by using solid ( ) and dashed molecule. In addition to these models,
computer graphics can also be used for
( ) wedge formula, the 3-D image of a
molecular modelling.
molecule from a two-dimensional picture
can be perceived. In these formulas the
solid-wedge is used to indicate a bond
projecting out of the plane of paper, towards
the observer. The dashed-wedge is used to
depict the bond projecting out of the plane of
the paper and away from the observer. Wedges
are shown in such a way that the broad end
of the wedge is towards the observer. The Ball and stick model
Framework model
bonds lying in plane of the paper are depicted
by using a normal line (—). 3-D representation
of methane molecule on paper has been
shown in Fig. 12.1.
Fig. 12.2
Fig. 12.1 Wedge-and-dash representation of CH4
ORGANIC CHEMISTRY – SOME BASIC PRINCIPLES AND TECHNIQUES 331
Cyclopropane Cyclohexane
Cyclohexene Tetrahydrofuran
These exhibit some of the properties similar
to those of aliphatic compounds.
Aromatic compounds
Aromatic compounds are special types of
compounds. You will learn about these
compounds in detail in Unit 13. These include
benzene and other related ring compounds
(benzenoid). Like alicyclic compounds,
aromatic comounds may also have hetero
atom in the ring. Such compounds are called
I. Acyclic or open chain compounds hetrocyclic aromatic compounds. Some of the
examples of various types of aromatic
These compounds are also called as aliphatic
compounds are:
compounds and consist of straight or
branched chain compounds, for example: Benzenoid aromatic compounds
CH3CH3
Ethane
Isobutane
Benzene Aniline Naphthalene
Non-benzenoid compound
Heterocyclic aromatic compounds acid found in red ant is named formic acid
since the Latin word for ant is formica. These
names are traditional and are considered as
trivial or common names. Some common
names are followed even today. For example,
Furan Thiophene Pyridine Buckminsterfullerene is a common name
Organic compounds can also be classified given to the newly discovered C60 cluster
on the basis of functional groups, into families (a form of carbon) noting its structural
or homologous series. similarity to the geodesic domes popularised
by the famous architect R. Buckminster
Functional Group
Fuller. Common names are useful and in
The functional group may be defined as an many cases indispensable, particularly when
atom or group of atoms joined in a specific the alternative systematic names are lengthy
manner which is responsible for the and complicated. Common names of some
characteristic chemical properties of the organic compounds are given in Table 12.1.
organic compounds. The examples are
Table 12.1 Common or Trivial Names of Some
hydroxyl group (–OH), aldehyde group (–CHO) Organic Compounds
and carboxylic acid group (–COOH) etc.
Homologous Series
A group or a series of organic compounds each
containing a characteristic functional group
forms a homologous series and the members
of the series are called homologues. The
members of a homologous series can be
represented by general molecular formula and
the successive members differ from each other
in molecular formula by a –CH2 unit. There
are a number of homologous series of
organic compounds. Some of these are
alkanes, alkenes, alkynes, haloalkanes,
alkanols, alkanals, alkanones, alkanoic acids,
amines etc.
12.5 NOMENCLATURE OF ORGANIC
COMPOUNDS
Organic chemistry deals with millions of
compounds. In order to clearly identify them, a
systematic method of naming has been 12.5.1 The IUPAC System of Nomenclature
developed and is known as the IUPAC A systematic name of an organic compound
(International Union of Pure and Applied is generally derived by identifying the parent
Chemistry) system of nomenclature. In this hydrocarbon and the functional group(s)
systematic nomenclature, the names are attached to it. See the example given below.
correlated with the structure such that the
reader or listener can deduce the structure from
the name.
Before the IUPAC system of nomenclature,
however, organic compounds were assigned
names based on their origin or certain
properties. For instance, citric acid is named
so because it is found in citrus fruits and the
ORGANIC CHEMISTRY – SOME BASIC PRINCIPLES AND TECHNIQUES 333
By further using prefixes and suffixes, the In order to name such compounds, the names
parent name can be modified to obtain the of alkyl groups are prefixed to the name of
actual name. Compounds containing carbon parent alkane. An alkyl group is derived from
and hydrogen only are called hydrocarbons. A a saturated hydrocarbon by removing a
hydrocarbon is termed saturated if it contains hydrogen atom from carbon. Thus, CH4
only carbon-carbon single bonds. The IUPAC becomes -CH3 and is called methyl group. An
name for a homologous series of such alkyl group is named by substituting ‘yl’ for
compounds is alkane. Paraffin (Latin: little ‘ane’ in the corresponding alkane. Some alkyl
affinity) was the earlier name given to these groups are listed in Table 12.3.
compounds. Unsaturated hydrocarbons are Table 12.3 Some Alkyl Groups
those, which contain at least one carbon-
carbon double or triple bond.
12.5.2 IUPAC Nomenclature of Alkanes
Straight chain hydrocarbons: The names
of such compounds are based on their chain
structure, and end with suffix ‘-ane’ and carry
a prefix indicating the number of carbon
atoms present in the chain (except from CH4
to C4H10, where the prefixes are derived from
trivial names). The IUPAC names of some
straight chain saturated hydrocarbons are
given in Table 12.2. The alkanes in Table 12.2 Abbreviations are used for some alkyl
differ from each other by merely the number groups. For example, methyl is abbreviated
of -CH 2 groups in the chain. They are as Me, ethyl as Et, propyl as Pr and butyl as
homologues of alkane series. Bu. The alkyl groups can be branched also.
Thus, propyl and butyl groups can have
Table 12.2 IUPAC Names of Some Unbranched branched structures as shown below.
Saturated Hydrocarbons
CH3-CH- CH3-CH2-CH- CH3-CH-CH2-
⏐ ⏐ ⏐
CH3 CH3 CH3
Isopropyl- sec-Butyl- Isobutyl-
CH3 CH3
⏐ ⏐
CH3-C- CH3-C-CH2-
⏐ ⏐
CH3 CH3
tert-Butyl- Neopentyl-
Common branched groups have specific
Branched chain hydrocarbons: In a
trivial names. For example, the propyl groups
branched chain compound small chains of
can either be n-propyl group or isopropyl
carbon atoms are attached at one or more
group. The branched butyl groups are called
carbon atoms of the parent chain. The small
sec-butyl, isobutyl and tert-butyl group. We
carbon chains (branches) are called alkyl
also encounter the structural unit,
groups. For example:
–CH2C(CH3)3, which is called neopentyl group.
CH3–CH–CH2–CH3 CH3–CH–CH2–CH–CH3
Nomenclature of branched chain alkanes:
⏐ ⏐ ⏐
We encounter a number of branched chain
CH3 CH2CH3 CH3
alkanes. The rules for naming them are given
(a) (b) below.
334 CHEMISTRY
1. First of all, the longest carbon chain in separated from the groups by hyphens and
the molecule is identified. In the example there is no break between methyl and
(I) given below, the longest chain has nine nonane.]
carbons and it is considered as the parent 4. If two or more identical substituent groups
or root chain. Selection of parent chain as are present then the numbers are
shown in (II) is not correct because it has separated by commas. The names of
only eight carbons. identical substituents are not repeated,
instead prefixes such as di (for 2), tri
(for 3), tetra (for 4), penta (for 5), hexa (for
6) etc. are used. While writing the name of
the substituents in alphabetical order,
these prefixes, however, are not considered.
Thus, the following compounds are
named as:
CH3 CH3 CH3 CH3
⏐ ⏐ ⏐ ⏐
CH3-CH-CH2-CH-CH3 CH3⎯C⎯CH2⎯CH⎯CH3
1 2 3 4 5 1 2⏐ 3 4 5
CH3
2. The carbon atoms of the parent chain are
numbered to identify the parent alkane and 2,4-Dimethylpentane 2,2,4-Trimethylpentane
to locate the positions of the carbon atoms H 3 C H2 C CH3
at which branching takes place due to the ⏐ ⏐
substitution of alkyl group in place of CH3⎯CH2⎯CH⎯C⎯CH2⎯CH2⎯CH3
hydrogen atoms. The numbering is done 1 2 3 ⏐4 5 6 7
in such a way that the branched carbon
CH3
atoms get the lowest possible numbers.
Thus, the numbering in the above example 3-Ethyl-4,4-dimethylheptane
should be from left to right (branching at
5. If the two substituents are found in
carbon atoms 2 and 6) and not from right
to left (giving numbers 4 and 8 to the equivalent positions, the lower number is
carbon atoms at which branches are given to the one coming first in the
attached). alphabetical listing. Thus, the following
compound is 3-ethyl-6-methyloctane and
1 2 3 4 5 6 7 8 9 not 6-ethyl-3-methyloctane.
C ⎯ C ⎯ C ⎯ C ⎯ C ⎯ C ⎯C ⎯ C ⎯ C
1 2 3 4 5 6 7 8
⏐ ⏐
CH3 — CH2—CH—CH2—CH2—CH—CH2 —CH3
C C⎯C
⏐ ⏐
9 8 7 6 5 4 3 2 1 CH2CH3 CH3
C⎯ C⎯C⎯C⎯C⎯C⎯C⎯C⎯C
⏐ ⏐ 6. The branched alkyl groups can be named
C C⎯C by following the above mentioned
3. The names of alkyl groups attached as a procedures. However, the carbon atom of
branch are then prefixed to the name of the branch that attaches to the root
the parent alkane and position of the alkane is numbered 1 as exemplified
substituents is indicated by the below.
appropriate numbers. If different alkyl 4 3 2 1
groups are present, they are listed in CH3–CH–CH2–CH–
alphabetical order. Thus, name for the ⏐ ⏐
compound shown above is: 6-ethyl-2- CH3 CH3
methylnonane. [Note: the numbers are 1,3-Dimethylbutyl-
ORGANIC CHEMISTRY – SOME BASIC PRINCIPLES AND TECHNIQUES 335
The name of such branched chain alkyl group Cyclic Compounds: A saturated monocyclic
is placed in parenthesis while naming the compound is named by prefixing ‘cyclo’ to the
compound. While writing the trivial names of corresponding straight chain alkane. If side
substituents’ in alphabetical order, the chains are present, then the rules given above
prefixes iso- and neo- are considered to be are applied. Names of some cyclic compounds
the part of the fundamental name of alkyl are given below.
group. The prefixes sec- and tert- are not
considered to be the part of the fundamental
name. The use of iso and related common
prefixes for naming alkyl groups is also
allowed by the IUPAC nomenclature as long
as these are not further substituted. In multi-
substituted compounds, the following rules
may aso be remembered:
• If there happens to be two chains of equal
size, then that chain is to be selected
which contains more number of side 3-Ethyl-1,1-dimethylcyclohexane
chains. (not 1-ethyl-3,3-dimethylcyclohexane)
• After selection of the chain, numbering is
to be done from the end closer to the Problem 12.7
substituent. Structures and IUPAC names of some
hydrocarbons are given below. Explain
why the names given in the parentheses
are incorrect.
2,5,6- Trimethyloctane
[and not 3,4,7-Trimethyloctane]
5-(2-Ethylbutyl)-3,3-dimethyldecane
[and not 5-(2,2-Dimethylbutyl)-3-ethyldecane]
3-Ethyl-5-methylheptane
[and not 5-Ethyl-3-methylheptane]
Solution
(a) Lowest locant number, 2,5,6 is lower
than 3,5,7, (b) substituents are in
5-sec-Butyl-4-isopropyldecane equivalent position; lower number is
given to the one that comes first in the
name according to alphabetical order.
chemical reactivity in an organic molecule. suffix. In such cases the full name of the parent
Compounds having the same functional group alkane is written before the class suffix. For
undergo similar reactions. For example, example CH 2 (OH)CH 2 (OH) is named as
CH3OH, CH3CH2OH, and (CH3)2CHOH — all ethane–1,2–diol. However, the ending – ne of
having -OH functional group liberate hydrogen the parent alkane is dropped in the case of
on reaction with sodium metal. The presence compounds having more than one double or
of functional groups enables systematisation triple bond; for example, CH2=CH-CH=CH2 is
of organic compounds into different classes. named as buta–1,3–diene.
Examples of some functional groups with their
prefixes and suf fixes along with some Problem 12.8
examples of organic compounds possessing Write the IUPAC names of the compounds
these are given in Table 12.4. i-iv from their given structures.
First of all, the functional group present
in the molecule is identified which determines
the choice of appropriate suffix. The longest
chain of carbon atoms containing the
functional group is numbered in such a way
that the functional group is attached at the Solution
carbon atom possessing lowest possible • The functional group present is an
number in the chain. By using the suffix as alcohol (OH). Hence the suffix is ‘-ol’.
given in Table 12.4, the name of the compound • The longest chain containing -OH has
is arrived at. eight carbon atoms. Hence the
In the case of polyfunctional compounds, corresponding saturated hydrocarbon
one of the functional groups is chosen as the is octane.
principal functional group and the compound is • The -OH is on carbon atom 3. In
then named on that basis. The remaining addition, a methyl group is attached
functional groups, which are subordinate at 6th carbon.
functional groups, are named as substituents Hence, the systematic name of this
using the appropriate prefixes. The choice of compound is 6-Methyloctan-3-ol.
principal functional group is made on the basis
of order of preference. The order of decreasing
priority for some functional groups is:
-COOH, –SO3H, -COOR (R=alkyl group), COCl,
-CONH2, -CN,-HC=O, >C=O, -OH, -NH2, >C=C<,
-C≡≡C- . Solution
The –R, C6H5-, halogens (F, Cl, Br, I), –NO2, The functional group present is ketone
alkoxy (–OR) etc. are always prefix (>C=O), hence suffix ‘-one’. Presence of
substituents. Thus, a compound containing two keto groups is indicated by ‘di’,
both an alcohol and a keto group is named hence suffix becomes ‘dione’. The two
as hydroxyalkanone since the keto group is keto groups are at carbons 2 and 4. The
preferred to the hydroxyl group. longest chain contains 6 carbon atoms,
For example, HOCH2(CH2)3CH2COCH3 will be hence, parent hydrocarbon is hexane.
named as 7-hydroxyheptan-2-one and not as Thus, the systematic name is Hexane-
2-oxoheptan -7-ol. Similarly, BrCH2CH=CH2 2,4-dione.
is named as 3-bromoprop-1-ene and not 1-
bromoprop-2-ene.
If more than one functional group of the
same type are present, their number is
indicated by adding di, tri, etc. before the class
ORGANIC CHEMISTRY – SOME BASIC PRINCIPLES AND TECHNIQUES 337
2-Chloro-1-methyl-4-nitrobenzene
(not 4-methyl-5-chloro-nitrobenzene)
CH3
⏐
CH3⎯ C⎯ CH3
(a) (b) ⏐
CH3
Neopentane
(2,2-Dimethylpropane)
with 1s orbital of each of the three hydrogen Alkyl radicals are classified as primary,
atoms. Each bond may be represented as secondary, or tertiary. Alkyl radical stability
C(sp 2)–H(1s) sigma bond. The remaining increases as we proceed from primary to
carbon orbital is perpendicular to the tertiary:
molecular plane and contains no electrons.
(Fig. 12.3). ,
Methyl Ethyl Isopropyl Tert-butyl
free free free free
radical radical radical radical
Organic reactions, which proceed by
homolytic fission are called free radical or
homopolar or nonpolar reactions.
12.7.2 Nucleophiles and Electrophiles
A reagent that brings an electron pair is called
Fig. 12.3 Shape of methyl cation a nucleophile (Nu:) i.e., nucleus seeking and
the reaction is then called nucleophilic. A
The heterolytic cleavage can also give a
reagent that takes away an electron pair is
species in which carbon gets the shared pair
called electrophile (E+) i.e., electron seeking
of electrons. For example, when group Z
and the reaction is called electrophilic.
attached to the carbon leaves without
During a polar organic reaction, a
nucleophile attacks an electrophilic centre of
the substrate which is that specific atom or
electron pair, the methyl anion is part of the electrophile that is electron
deficient. Similarly, the electrophiles attack at
formed. Such a carbon species carrying a nucleophilic centre, which is the electron
negative charge on carbon atom is called rich centre of the substrate. Thus, the
carbanion. Carbanions are also unstable and electrophiles receive electron pair from
reactive species. The organic reactions which nucleophile when the two undergo bonding
proceed through heterolytic bond cleavage are interaction. A curved-arrow notation is used
called ionic or heteropolar or just polar to show the movement of an electron pair from
reactions. the nucleophile to the electrophile. Some
In homolytic cleavage, one of the examples of nucleophiles are the negatively
electrons of the shared pair in a covalent bond charged ions with lone pair of electrons such
– –
goes with each of the bonded atoms. Thus, in as hydroxide (HO ), cyanide (NC ) ions and
–
homolytic cleavage, the movement of a single carbanions (R3C: ). Neutral molecules such
electron takes place instead of an electron
as etc., can also act as
pair. The single electron movement is shown
nucleophiles due to the presence of lone pair
by ‘half-headed’ (fish hook: ) curved arrow.
of electrons. Examples of electrophiles
Such cleavage results in the formation of +
neutral species (atom or group) which include carbocations ( C H 3 ) and neutral
contains an unpaired electron. These species molecules having functional groups like
are called free radicals. Like carbocations carbonyl group (>C=O) or alkyl halides
and carbanions, free radicals are also (R 3C-X, where X is a halogen atom). The
very reactive. A homolytic cleavage can be carbon atom in carbocations has sextet
shown as: configuration; hence, it is electron deficient
and can receive a pair of electrons from the
nucleophiles. In neutral molecules such as
Alkyl alkyl halides, due to the polarity of the C-X
free radical bond a partial positive charge is generated
ORGANIC CHEMISTRY – SOME BASIC PRINCIPLES AND TECHNIQUES 343
on the carbon atom and hence the carbon 12.7.3 Electron Movement in Organic
atom becomes an electrophilic centre at Reactions
which a nucleophile can attack. The movement of electrons in organic
reactions can be shown by curved-arrow
Problem 12.11 notation. It shows how changes in bonding
Using curved-arrow notation, show the occur due to electronic redistribution during
formation of reactive intermediates when the reaction. To show the change in position
the following covalent bonds undergo of a pair of electrons, curved arrow starts from
heterolytic cleavage. the point from where an electron pair is shifted
(a) CH3–SCH3, (b) CH3–CN, (c) CH3–Cu and it ends at a location to which the pair of
electron may move.
Solution
Presentation of shifting of electron pair is
given below :
when a reagent approaches to attack it. This nitro (- NO2), cyano (- CN), carboxy (- COOH),
type of electron displacement is called ester (-COOR), aryloxy (-OAr, e.g. – OC6H5),
electromeric effect or polarisability effect. In etc. are electron-withdrawing groups. On the
the following sections we will learn about these other hand, the alkyl groups like methyl
types of electronic displacements. (–CH 3) and ethyl (–CH 2–CH 3) are usually
considered as electron donating groups.
12.7.5 Inductive Effect
When a covalent bond is formed between Problem 12.14
atoms of different electronegativity, the Which bond is more polar in the following
electron density is more towards the more pairs of molecules: (a) H3C-H, H3C-Br
electronegative atom of the bond. Such a shift (b) H 3 C-NH 2 , H 3 C-OH (c) H 3 C-OH,
of electron density results in a polar covalent H3C-SH
bond. Bond polarity leads to various electronic
effects in organic compounds. Solution
Let us consider cholorethane (CH3CH2Cl) (a) C–Br, since Br is more electronegative
in which the C–Cl bond is a polar covalent than H, (b) C–O, (c) C–O
bond. It is polarised in such a way that the Problem 12.15
+
carbon-1 gains some positive charge (δ ) and
– In which C–C bond of CH3CH2CH2Br, the
the chlorine some negative charge (δ ). The
fractional electronic charges on the two atoms inductive effect is expected to be the
in a polar covalent bond are denoted by least?
symbol δ (delta) and the shift of electron Solution
density is shown by an arrow that points from
+ – Magnitude of inductive effect diminishes
δ to δ end of the polar bond.
+ + − as the number of intervening bonds
δδ δ δ
increases. Hence, the effect is least in the
CH3 ⎯→⎯CH2⎯→⎯ ⎯→⎯Cl
bond between carbon-3 and hydrogen.
2 1
In turn carbon-1, which has developed 12.7.6 Resonance Structure
+
partial positive charge (δ ) draws some
There are many organic molecules whose
electron density towards it from the adjacent
behaviour cannot be explained by a single
C-C bond. Consequently, some positive charge
+ + Lewis structure. An example is that of
(δδ ) develops on carbon-2 also, where δδ
benzene. Its cyclic structure
symbolises relatively smaller positive charge
containing alternating C–C single
as compared to that on carbon – 1. In other
words, the polar C – Cl bond induces polarity and C=C double bonds shown is
in the adjacent bonds. Such polarisation of inadequate for explaining its Benzene
σ-bond caused by the polarisation of adjacent characteristic properties.
σ-bond is referred to as the inductive effect. As per the above representation, benzene
This effect is passed on to the subsequent should exhibit two different bond lengths, due
bonds also but the effect decreases rapidly to C–C single and C=C double bonds. However,
as the number of intervening bonds increases as determined experimentally benzene has a
and becomes vanishingly small after three uniform C–C bond distances of 139 pm, a
bonds. The inductive effect is related to the value inter mediate between the C–C
ability of substituent(s) to either withdraw or single(154 pm) and C=C double (134 pm)
donate electron density to the attached carbon bonds. Thus, the structure of benzene cannot
atom. Based on this ability, the substitutents be represented adequately by the above
can be classified as electron-withdrawing or structure. Further, benzene can be
electron donating groups relative to hydrogen. represented equally well by the energetically
Halogens and many other groups such as identical structures I and II.
ORGANIC CHEMISTRY – SOME BASIC PRINCIPLES AND TECHNIQUES 345
Problem 12.17
Write resonance structures of
CH2=CH–CHO. Indicate relative stability of
However, it is known that the two N–O the contributing structures.
bonds of nitromethane are of the same
length (intermediate between a N–O single Solution
bond and a N=O double bond). The actual
structure of nitromethane is therefore a
resonance hybrid of the two canonical
forms I and II.
The energy of actual structure of the
molecule (the resonance hybrid) is lower than
that of any of the canonical structures. The
difference in energy between the actual
structure and the lowest energy resonance
structure is called the resonance Stability: I > II > III
stabilisation energy or simply the [I: Most stable, more number of covalent
resonance energy. The more the number of bonds, each carbon and oxygen atom has
important contributing structures, the more an octet and no separation of opposite
is the resonance energy. Resonance is charge II: negative charge on more
particularly important when the contributing electronegative atom and positive charge
structures are equivalent in energy. on more electropositive atom; III: does
The following rules are applied while writing not contribute as oxygen has positive
resonance structures: charge and carbon has negative charge,
The resonance structures have (i) the same hence least stable].
positions of nuclei and (ii) the same number of
346 CHEMISTRY
When inductive and electromeric effects In general, greater the number of alkyl
operate in opposite directions, the electomeric groups attached to a positively charged carbon
effect predominates. atom, the greater is the hyperconjugation
interaction and stabilisation of the cation.
12.7.9 Hyperconjugation
Thus, we have the following relative stability
Hyperconjugation is a general stabilising of carbocations :
interaction. It involves delocalisation of
σ electrons of C—H bond of an alkyl group
directly attached to an atom of unsaturated
system or to an atom with an unshared
p orbital. The σ electrons of C—H bond of the
alkyl group enter into partial conjugation with Hyperconjugation is also possible in
the attached unsaturated system or with the alkenes and alkylarenes.
unshared p orbital. Hyperconjugation is a Delocalisation of electrons by
permanent effect. hyperconjugation in the case of alkene can
To understand hyperconjugation effect, let be depicted as in Fig. 12.4(b).
+
us take an example of CH3 CH2 (ethyl cation)
in which the positively charged carbon atom
has an empty p orbital. One of the C-H bonds
of the methyl group can align in the plane of
this empty p orbital and the electrons
constituting the C-H bond in plane with this
p orbital can then be delocalised into the
empty p orbital as depicted in Fig. 12.4 (a).
Fig. 12.4(b) Orbital diagram showing
hyperconjugation in propene
There are various ways of looking at the
hyperconjugative effect. One of the way is to
regard C—H bond as possessing partial ionic
character due to resonance.
Fig.12.6 Fractional distillation. The vapours of lower boiling fraction reach the
top of the column first followed by vapours of higher boiling fractions.
350 CHEMISTRY
Vapours of the liquid with higher boiling theoretical plate. Commercially, columns
point condense before the vapours of the with hundreds of plates are available.
liquid with lower boiling point. The vapours One of the technological applications of
rising up in the fractionating column become fractional distillation is to separate different
richer in more volatile component. By the time fractions of crude oil in petroleum industry.
the vapours reach to the top of the
Distillation under reduced pressure: This
fractionating column, these are rich in the
method is used to purify liquids having very
more volatile component. Fractionating
high boiling points and those, which
columns are available in various sizes and
decompose at or below their boiling points.
designs as shown in Fig.12.7. A fractionating
Such liquids are made to boil at a temperature
column provides many surfaces for heat
lower than their normal boiling points by
exchange between the ascending vapours
reducing the pressure on their surface. A
and the descending condensed liquid. Some
liquid boils at a temperature at which its
of the condensing liquid in the fractionating
vapour pressure is equal to the external
column obtains heat from the ascending
pressure. The pressure is reduced with the
vapours and revaporises. The vapours thus
help of a water pump or vacuum pump
become richer in low boiling component. The
(Fig.12.8). Glycerol can be separated from
vapours of low boiling component ascend to
spent-lye in soap industry by using this
the top of the column. On reaching the top,
technique.
the vapours become pure in low boiling
component and pass through the condenser Steam Distillation: This technique is
and the pure liquid is collected in a receiver. applied to separate substances which are
After a series of successive distillations, the steam volatile and are immiscible with
remaining liquid in the distillation flask gets water. In steam distillation, steam from a
enriched in high boiling component. Each steam generator is passed through a heated
successive condensation and vaporisation flask containing the liquid to be distilled.
unit in the fractionating column is called a The mixture of steam and the volatile
organic compound is condensed and
collected. The compound is later separated
from water using a separating funnel. In
steam distillation, the liquid boils when
the sum of vapour pressures due to the
organic liquid (p 1 ) and that due to water
(p 2 ) becomes equal to the atmospheric
pressure (p), i.e. p =p 1 + p 2 . Since p 1 is
lower than p, the organic liquid vaporises
at lower temperature than its boiling
point.
Thus, if one of the substances in the
mixture is water and the other, a water
insoluble substance, then the mixture will boil
close to but below, 373K. A mixture of water
and the substance is obtained which can be
separated by using a separating funnel.
Aniline is separated by this technique from
aniline – water mixture (Fig.12.9).
12.8.4 Differential Extraction
When an organic compound is present in an
Fig.12.7 Different types of fractionating columns. aqueous medium, it is separated by shaking
ORGANIC CHEMISTRY – SOME BASIC PRINCIPLES AND TECHNIQUES 351
Fig.12.8 Distillation under reduced pressure. A liquid boils at a temperature below its
vapour pressure by reducing the pressure.
Fig.12.9 Steam distillation. Steam volatile component volatilizes, the vapours condense in
the condenser and the liquid collects in conical flask.
352 CHEMISTRY
it with an organic solvent in which it is more mixture get gradually separated from one
soluble than in water. The organic solvent and another. The moving phase is called the mobile
the aqueous solution should be immiscible phase.
with each other so that they form two distinct Based on the principle involved,
layers which can be separated by separatory chromatography is classified into different
funnel. The organic solvent is later removed categories. Two of these are:
by distillation or by evaporation to get back (a) Adsorption chromatography, and
the compound. Differential extraction is (b) Partition chromatography.
carried out in a separatory funnel as shown
in Fig. 12.10. If the organic compound is less a) Adsorption Chromatography: Adsor -
ption chromatography is based on the fact
that different compounds are adsorbed on an
adsorbent to different degrees. Commonly
used adsorbents are silica gel and alumina.
When a mobile phase is allowed to move
over a stationary phase (adsorbent),
the components of the mixture move by
varying distances over the stationary
phase. Following are two main types of
chromatographic techniques based on the
principle of differential adsorption.
(a) Column chromatography, and
(b) Thin layer chromatography.
Column Chromatography: Column
chromatography involves separation of a
Fig.12.10 Differential extraction. Extraction of com- mixture over a column of adsorbent
pound takes place based on difference (stationary phase) packed in a glass tube. The
in solubility
column is fitted with a stopcock at its lower
soluble in the organic solvent, a very large end (Fig. 12.11). The mixture adsorbed on
quantity of solvent would be required to
extract even a very small quantity of the
compound. The technique of continuous
extraction is employed in such cases. In this
technique same solvent is repeatedly used for
extraction of the compound.
12.8.5 Chromatography
Chromatography is an important technique
extensively used to separate mixtures into
their components, purify compounds and also
to test the purity of compounds. The name
chromatography is based on the Greek word
chroma, for colour since the method was first
used for the separation of coloured
substances found in plants. In this technique,
the mixture of substances is applied onto a
stationary phase, which may be a solid or a
liquid. A pure solvent, a mixture of solvents, Fig.12.11 Column chromatography. Different
or a gas is allowed to move slowly over the stages of separation of components
stationary phase. The components of the of a mixture.
ORGANIC CHEMISTRY – SOME BASIC PRINCIPLES AND TECHNIQUES 353
adsorbent is placed on the top of the eluant rises up the plate, the components of
adsorbent column packed in a glass tube. An the mixture move up along with the eluant to
appropriate eluant which is a liquid or a different distances depending on their degree
mixture of liquids is allowed to flow down the of adsorption and separation takes place. The
column slowly. Depending upon the degree relative adsorption of each component of the
to which the compounds are adsorbed, mixture is expressed in ter ms of its
complete separation takes place. The most retardation factor i.e. Rf value (Fig.12.12 b).
readily adsorbed substances are retained near Distance moved by the substance from base line (x)
the top and others come down to various Rf =
Distance moved by the solvent from base line (y)
distances in the column (Fig.12.11).
The spots of coloured compounds are
Thin Layer Chromatography: Thin layer
visible on TLC plate due to their original
chromatography (TLC) is another type of
colour. The spots of colourless compounds,
adsorption chromatography, which involves
which are invisible to the eye but fluoresce,
separation of substances of a mixture over a
can be detected by putting the plate under
thin layer of an adsorbent coated on glass
plate. A thin layer (about 0.2mm thick) of an ultraviolet light. Another detection technique
adsorbent (silica gel or alumina) is spread over is to place the plate in a covered jar containing
a glass plate of suitable size. The plate is a few crystals of iodine. Spots of compounds,
known as thin layer chromatography plate or which adsorb iodine, will show up as brown
chromaplate. The solution of the mixture to spots. Sometimes an appropriate reagent may
be separated is applied as a small spot about also be sprayed on the plate. For example,
2 cm above one end of the TLC plate. The amino acids may be detected by spraying the
glass plate is then placed in a closed jar plate with ninhydrin solution (Fig.12.12b).
containing the eluant (Fig. 12.12a). As the Partition Chromatography: Partition
chromatography is based on continuous
differential partitioning of components of a
mixture between stationary and mobile
phases. Paper chromatography is a type of
partition chromatography. In paper
chromatography, a special quality paper
known as chromatography paper is used.
Chromatography paper contains water
trapped in it, which acts as the stationary
phase.
A strip of chromatography paper spotted
Fig.12.12 (a) Thin layer chromatography. at the base with the solution of the mixture is
Chromatogram being developed. suspended in a suitable solvent or a mixture
of solvents (Fig. 12.13). This solvent acts as
the mobile phase. The solvent rises up the
paper by capillary action and flows over the
spot. The paper selectively retains different
components according to their differing
partition in the two phases. The paper strip
so developed is known as a chromatogram.
The spots of the separated coloured
compounds are visible at different heights
from the position of initial spot on the
chromatogram. The spots of the separated
Fig.12.12 (b) Developed chromatogram. colourless compounds may be observed either
354 CHEMISTRY
Fig.12.14 Estimation of carbon and hydrogen. Water and carbon dioxide formed on oxidation of substance
are absorbed in anhydrous calcium chloride and potassium hydroxide solutions respectively
contained in U tubes.
356 CHEMISTRY
Fig.12.15 Dumas method. The organic compound yields nitrogen gas on heating it with
copper(II) oxide in the presence of CO2 gas. The mixture of gases is collected
over potassium hydroxide solution in which CO2 is absorbed and volume of
nitrogen gas is determined.
ORGANIC CHEMISTRY – SOME BASIC PRINCIPLES AND TECHNIQUES 357
Fig.12.16 Kjeldahl method. Nitrogen-containing compound is treated with concentrated H2SO4 to get
ammonium sulphate which liberates ammonia on treating with NaOH; ammonia is absorbed
in known volume of standard acid.
358 CHEMISTRY
present in the compound are oxidised to 1 mol of BaSO4 = 233 g BaSO4 = 32 g sulphur
carbon dioxide and water. The halogen
32 × m1
present forms the corresponding silver halide m1 g BaSO4 contains g sulphur
(AgX). It is filtered, washed, dried and weighed. 233
Let the mass of organic 32 × m1 × 100
compound taken = m g Percentage of sulphur=
Mass of AgX formed = m1 g 233 × m
1 mol of AgX contains 1 mol of X
Problem 12.24
Mass of halogen in m1g of AgX
In sulphur estimation, 0.157 g of an
atomic mass of X × m1g organic compound gave 0.4813 g of
=
molecular mass of AgX barium sulphate. What is the
Percentage of halogen percentage of sulphur in the compound?
Solution
atomic mass of X × m1 × 100
= Molecular mass of BaSO4 = 137+32+64
molecular mass of AgX × m = 233 g
Problem 12.23 233 g BaSO4 contains 32 g sulphur
In Carius method of estimation of 32 × 0.4813
halogen, 0.15 g of an organic compound 0.4813 g BaSO4 contains g
233
gave 0.12 g of AgBr. Find out the
sulphur
percentage of bromine in the compound.
Solution 32 × 0.4813 × 100
Percentage of sulphur=
Molar mass of AgBr = 108 + 80 233 × 0.157
= 188 g mol-1 = 42.10%
188 g AgBr contains 80 g bromine
12.10.5 Phosphorus
80 × 0.12
0.12 g AgBr contains g bromine A known mass of an organic compound is
188 heated with fuming nitric acid whereupon
phosphorus present in the compound is
80 × 0.12 × 100
Percentage of bromine= oxidised to phosphoric acid. It is precipitated
188 × 0.15 as ammonium phosphomolybdate, (NH4) 3
= 34.04% PO 4 .12MoO 3 , by adding ammonia and
ammonium molybdate. Alter natively,
12.10.4 Sulphur phosphoric acid may be precipitated as
A known mass of an organic compound is MgNH 4 PO 4 by adding magnesia mixture
heated in a Carius tube with sodium peroxide which on ignition yields Mg2P2O7.
or fuming nitric acid. Sulphur present in the Let the mass of organic compound taken
compound is oxidised to sulphuric acid. It is = m g and mass of ammonium phospho
precipitated as barium sulphate by adding molydate = m1g
excess of barium chloride solution in water.
Molar mass of (NH4)3PO4.12MoO3 = 1877 g
The precipitate is filtered, washed, dried and
weighed. The percentage of sulphur can be 31 × m1 × 100
calculated from the mass of barium sulphate. Percentage of phosphorus = %
1877 × m
Let the mass of organic
If phosphorus is estimated as Mg2P2O7,
compound taken = m g
and the mass of barium 62 × m1 × 100
sulphate formed = m1g Percentage of phosphorus = %
222 × m
360 CHEMISTRY
SUMMARY
In this unit, we have learnt some basic concepts in structure and reactivity of organic
compounds, which are formed due to covalent bonding. The nature of the covalent bonding
in organic compounds can be described in terms of orbitals hybridisation concept, according
to which carbon can have sp3, sp2 and sp hybridised orbitals. The sp3, sp2 and sp hybridised
carbons are found in compounds like methane, ethene and ethyne respectively. The
tetrahedral shape of methane, planar shape of ethene and linear shape of ethyne can be
understood on the basis of this concept. A sp3 hybrid orbital can overlap with 1s orbital of
hydrogen to give a carbon - hydrogen (C–H) single bond (sigma, σ bond). Overlap of a sp2
orbital of one carbon with sp2 orbital of another results in the formation of a carbon–carbon
σ bond. The unhybridised p orbitals on two adjacent carbons can undergo lateral (side-by-
side) overlap to give a pi (π) bond. Organic compounds can be represented by various structural
formulas. The three dimensional representation of organic compounds on paper can be
drawn by wedge and dash formula.
Organic compounds can be classified on the basis of their structure or the functional
groups they contain. A functional group is an atom or group of atoms bonded together in a
unique fashion and which determines the physical and chemical properties of the compounds.
The naming of the organic compounds is carried out by following a set of rules laid down by
the International Union of Pure and Applied Chemistry (IUPAC). In IUPAC nomenclature,
the names are correlated with the structure in such a way that the reader can deduce the
structure from the name.
ORGANIC CHEMISTRY – SOME BASIC PRINCIPLES AND TECHNIQUES 361
Organic reaction mechanism concepts are based on the structure of the substrate
molecule, fission of a covalent bond, the attacking reagents, the electron displacement effects
and the conditions of the reaction. These organic reactions involve breaking and making of
covalent bonds. A covalent bond may be cleaved in heterolytic or homolytic fashion. A
heterolytic cleavage yields carbocations or carbanions, while a homolytic cleavage gives
free radicals as reactive intermediate. Reactions proceeding through heterolytic cleavage
involve the complimentary pairs of reactive species. These are electron pair donor known as
nucleophile and an electron pair acceptor known as electrophile. The inductive, resonance,
electromeric and hyperconjugation effects may help in the polarisation of a bond making
certain carbon atom or other atom positions as places of low or high electron densities.
Organic reactions can be broadly classified into following types; substitution, addition,
elimination and rearrangement reactions.
Purification, qualitative and quantitative analysis of organic compounds are carried out
for determining their structures. The methods of purification namely : sublimation, distillation
and differential extraction are based on the difference in one or more physical properties.
Chromatography is a useful technique of separation, identification and purification of
compounds. It is classified into two categories : adsorption and partition chromatography.
Adsorption chromatography is based on differential adsorption of various components of a
mixture on an adsorbent. Partition chromatography involves continuous partitioning of the
components of a mixture between stationary and mobile phases. After getting the compound
in a pure form, its qualitative analysis is carried out for detection of elements present in it.
Nitrogen, sulphur, halogens and phosphorus are detected by Lassaigne’s test. Carbon and
hydrogen are estimated by determining the amounts of carbon dioxide and water produced.
Nitrogen is estimated by Dumas or Kjeldahl’s method and halogens by Carius method.
Sulphur and phosphorus are estimated by oxidising them to sulphuric and phosphoric
acids respectively. The percentage of oxygen is usually determined by difference between
the total percentage (100) and the sum of percentages of all other elements present.
EXERCISES
12.1 What are hybridisation states of each carbon atom in the following compounds ?
CH2=C=O, CH3CH=CH2, (CH3)2CO, CH2=CHCN, C6H6
12.2 Indicate the σ and π bonds in the following molecules :
C6H6, C6H12, CH2Cl2, CH2=C=CH2, CH3NO2, HCONHCH3
12.3 Write bond line formulas for : Isopropyl alcohol, 2,3-Dimethyl butanal, Heptan-4-
one.
12.4 Give the IUPAC names of the following compounds :
12.5 Which of the following represents the correct IUPAC name for the compounds
concer ned ? (a) 2,2-Dimethylpentane or 2-Dimethylpentane (b) 2,4,7-
Trimethyloctane or 2,5,7-Trimethyloctane (c) 2-Chloro-4-methylpentane or
4-Chloro-2-methylpentane (d) But-3-yn-1-ol or But-4-ol-1-yne.
362 CHEMISTRY
12.6 Draw formulas for the first five members of each homologous series beginning with
the following compounds. (a) H–COOH (b) CH3COCH3 (c) H–CH=CH2
12.7 Give condensed and bond line structural formulas and identify the functional
group(s) present, if any, for :
(a) 2,2,4-Trimethylpentane
(b) 2-Hydroxy-1,2,3-propanetricarboxylic acid
(c) Hexanedial
12.8 Identify the functional groups in the following compounds
12.9 Which of the two: O2NCH2CH2O– or CH3CH2O– is expected to be more stable and
why ?
12.10 Explain why alkyl groups act as electron donors when attached to a π system.
12.11 Draw the resonance structures for the following compounds. Show the electron
shift using curved-arrow notation.
+
(a) C 6 H 5 OH (b) C 6 H 5NO 2 (c) CH 3 CH=CHCHO (d) C 6 H 5 –CHO (e) C6H5 −CH2
+
(f) CH3CH = CHC H2
–
(b) CH3COCH3 + C N → ( CH3 )2 C ( CN )( OH )
+
(c) C6H5 + CH3 C O → C6H5COCH3
12.14 Classify the following reactions in one of the reaction type studied in this unit.
(c) CH 3CH 2 Br + HO − → CH 2 = CH 2 + H 2O
(a)
ORGANIC CHEMISTRY – SOME BASIC PRINCIPLES AND TECHNIQUES 363
(b)
(c)
12.16 For the following bond cleavages, use curved-arrows to show the electron flow
and classify each as homolysis or heterolysis. Identify reactive intermediate
produced as free radical, carbocation and carbanion.
(a)
(b)
(c)
(d)
12.17 Explain the terms Inductive and Electromeric effects. Which electron displacement
effect explains the following correct orders of acidity of the carboxylic acids?
(a) Cl3CCOOH > Cl2CHCOOH > ClCH2COOH
(b) CH3CH2COOH > (CH3)2CHCOOH > (CH3)3C.COOH
12.18 Give a brief description of the principles of the following techniques taking an
example in each case.
(a) Crystallisation (b) Distillation (c) Chromatography
12.19 Describe the method, which can be used to separate two compounds with different
solubilities in a solvent S.
12.20 What is the difference between distillation, distillation under reduced pressure
and steam distillation ?
12.21 Discuss the chemistry of Lassaigne’s test.
12.22 Differentiate between the principle of estimation of nitrogen in an organic compound
by (i) Dumas method and (ii) Kjeldahl’s method.
12.23 Discuss the principle of estimation of halogens, sulphur and phosphorus present
in an organic compound.
12.24 Explain the principle of paper chromatography.
12.25 Why is nitric acid added to sodium extract before adding silver nitrate for testing
halogens?
12.26 Explain the reason for the fusion of an organic compound with metallic sodium
for testing nitrogen, sulphur and halogens.
12.27 Name a suitable technique of separation of the components from a mixture of
calcium sulphate and camphor.
12.28 Explain, why an organic liquid vaporises at a temperature below its boiling point
in its steam distillation ?
12.29 Will CCl4 give white precipitate of AgCl on heating it with silver nitrate? Give
reason for your answer.
364 CHEMISTRY
12.30 Why is a solution of potassium hydroxide used to absorb carbon dioxide evolved
during the estimation of carbon present in an organic compound?
12.31 Why is it necessary to use acetic acid and not sulphuric acid for acidification of
sodium extract for testing sulphur by lead acetate test?
12.32 An organic compound contains 69% carbon and 4.8% hydrogen, the remainder
being oxygen. Calculate the masses of carbon dioxide and water produced when
0.20 g of this substance is subjected to complete combustion.
12.33 A sample of 0.50 g of an organic compound was treated according to Kjeldahl’s
method. The ammonia evolved was absorbed in 50 ml of 0.5 M H2SO4. The residual
acid required 60 mL of 0.5 M solution of NaOH for neutralisation. Find the
percentage composition of nitrogen in the compound.
12.34 0.3780 g of an organic chloro compound gave 0.5740 g of silver chloride in Carius
estimation. Calculate the percentage of chlorine present in the compound.
12.35 In the estimation of sulphur by Carius method, 0.468 g of an organic sulphur
compound afforded 0.668 g of barium sulphate. Find out the percentage of sulphur
in the given compound.
12.36 In the organic compound CH2 = CH – CH2 – CH2 – C ≡ CH, the pair of hydridised
orbitals involved in the formation of: C2 – C3 bond is:
(a) sp – sp2 (b) sp – sp3 (c) sp2 – sp3 (d) sp3 – sp3
12.37 In the Lassaigne’s test for nitrogen in an organic compound, the Prussian blue
colour is obtained due to the formation of:
(a) Na4[Fe(CN)6] (b) Fe4[Fe(CN)6]3 (c) Fe2[Fe(CN)6] (d) Fe3[Fe(CN)6]4
12.38 Which of the following carbocation is most stable ?
+ + + +
(a) (CH3)3C. C H2 (b) (CH3)3 C (c) CH3CH2 C H2 (d) CH3 C H CH2CH3
12.39 The best and latest technique for isolation, purification and separation of organic
compounds is:
(a) Crystallisation (b) Distillation (c) Sublimation (d) Chromatography
12.40 The reaction:
CH3CH2I + KOH(aq) → CH3CH2OH + KI
is classified as :
(a) electrophilic substitution (b) nucleophilic substitution
(c) elimination (d) addition
HYDROCARBONS 365
UNIT 13
HYDROCARBONS
(ii) unsaturated and (iii) aromatic general formula for alkane family or
hydrocarbons. Saturated hydrocarbons homologous series? The general formula for
contain carbon-carbon and carbon-hydrogen alkanes is CnH2n+2, where n stands for number
single bonds. If different carbon atoms are of carbon atoms and 2n+2 for number of
joined together to form open chain of carbon hydrogen atoms in the molecule. Can you
atoms with single bonds, they are termed as recall the structure of methane? According to
alkanes as you have already studied in VSEPR theory (Unit 4), methane has a
Unit 12. On the other hand, if carbon atoms tetrahedral structure (Fig. 13.1) which is
form a closed chain or a ring, they are termed multiplanar, in which carbon atom lies at the
as cycloalkanes. Unsaturated hydrocarbons centre and the four hydrogen atoms lie at the
contain carbon-carbon multiple bonds – four corners of a regular tetrahedron. All
double bonds, triple bonds or both. Aromatic H-C-H bond angles are of 109.5°.
hydrocarbons are a special type of cyclic
compounds. You can construct a large number
of models of such molecules of both types
(open chain and close chain) keeping in mind
that carbon is tetravalent and hydrogen is
monovalent. For making models of alkanes,
you can use toothpicks for bonds and
plasticine balls for atoms. For alkenes, alkynes
and aromatic hydrocarbons, spring models can
be constructed. Fig. 13.1 Structure of methane
In alkanes, tetrahedra are joined together
13.2 ALKANES in which C-C and C-H bond lengths are
As already mentioned, alkanes are saturated 154 pm and 112 pm respectively (Unit 12). You
open chain hydrocarbons containing have already read that C–C and C–H σ bonds
3
carbon - carbon single bonds. Methane (CH4) are formed by head-on overlapping of sp
is the first member of this family. Methane is a hybrid orbitals of carbon and 1s orbitals of
gas found in coal mines and marshy places. If hydrogen atoms.
you replace one hydrogen atom of methane by
13.2.1 Nomenclature and Isomerism
carbon and join the required number of
hydrogens to satisfy the tetravalence of the You have already read about nomenclature
other carbon atom, what do you get? You get of different classes of organic compounds in
C 2 H 6 . This hydrocarbon with molecular Unit 12. Nomenclature and isomerism in
formula C2H6 is known as ethane. Thus you alkanes can further be understood with the
can consider C2H6 as derived from CH4 by help of a few more examples. Common names
replacing one hydrogen atom by -CH3 group. are given in parenthesis. First three alkanes
Go on constructing alkanes by doing this – methane, ethane and propane have only
theoretical exercise i.e., replacing hydrogen one structure but higher alkanes can have
atom by –CH3 group. The next molecules will more than one structure. Let us write
be C3H8, C4H10 … structures for C4H10. Four carbon atoms of
C4H10 can be joined either in a continuous
chain or with a branched chain in the
following two ways :
I
These hydrocarbons are inert under
normal conditions as they do not react with
acids, bases and other reagents. Hence, they
were earlier known as paraffins (latin : parum,
little; affinis, affinity). Can you think of the Butane (n- butane), (b.p. 273 K)
HYDROCARBONS 367
Solution
(i) CH3 – CH2 – CH2 – CH2– CH2– CH3
n-Hexane
Pentane (n-pentane)
(b.p. 309 K)
IV
2-Methylpentane
3-Methylpentane
2-Methylbutane (isopentane)
(b.p. 301 K)
2,3-Dimethylbutane
V
2,2 - Dimethylbutane
Problem 13.2
Write structures of different isomeric alkyl groups corresponding to the molecular formula
C5H11. Write IUPAC names of alcohols obtained by attachment of –OH groups at different
carbons of the chain.
Solution
Structures of – C5H11 group Corresponding alcohols
(i) CH3 – CH2 – CH2 – CH2– CH2 – CH3 – CH2 – CH2 – CH2– CH2 – OH
Pentan-1-ol
(ii) CH3 – CH – CH2 – CH2 – CH3 CH3 – CH – CH2 – CH2– CH3
| |
OH
Pentan-2-ol
(iii) CH3 – CH2 – CH – CH2 – CH2 CH3 – CH2 – CH – CH2– CH3
| |
OH
Pentan-3-ol
CH3 CH3
| |
(iv) CH3 – CH – CH2 – CH2 – CH3 – CH – CH2 – CH2– OH
3-Methylbutan-1-ol
CH3 CH3
| |
(v) CH3 – CH2 – CH – CH2 – CH3 – CH2 – CH – CH2– OH
2-Methylbutan-1-ol
CH3 CH3
| |
(vi) CH3 – C – CH2 – CH3 CH3 – C – CH2 – CH3
| |
OH
2-Methylbutan-2-ol
CH3 CH3
| |
(vii) CH3 – C – CH2 – CH3 – C – CH2OH
| |
CH3 CH3
2,2- Dimethylpropan-1-ol
HYDROCARBONS 369
(3,3-Diethyl-5-isopropyl-4-methyloctane)
5-(2,2– Dimethylpropyl)nonane
1
(e) CH3 – 2CH2 – 3CH – 4CH2 – 5CH – 6CH2 – 7CH3
Alphabetical
priority order
3–Ethyl–5–methylheptane
iii) Attach ethyl group at carbon 3 and two Longest chain is of six carbon atoms and
methyl groups at carbon 2 not that of five. Hence, correct name is
CH3 3-Methylhexane.
| 7 6 5 4 3 2 1
1 2 3 4 5
C – C– C– C– C (ii) CH3 – CH2 – CH – CH2 – CH – CH2 – CH3
| |
CH3 C2 H5
iv) Satisfy the valence of each carbon atom by
putting requisite number of hydrogen Numbering is to be started from the end
atoms : which gives lower number to ethyl group.
Hence, correct name is 3-ethyl-5-
CH3
| methylheptane.
CH3 – C – CH – CH2 – CH3
| | 13.2.2 Preparation
CH3 C2H5 Petroleum and natural gas are the main
Thus we arrive at the correct structure. If sources of alkanes. However, alkanes can be
you have understood writing of structure from prepared by following methods :
the given name, attempt the following
problems. 1. From unsaturated hydrocarbons
Dihydrogen gas adds to alkenes and alkynes
Problem 13.4 in the presence of finely divided catalysts like
Write structural formulas of the following platinum, palladium or nickel to form alkanes.
compounds : This process is called hydrogenation. These
metals adsorb dihydrogen gas on their surfaces
(i) 3, 4, 4, 5–Tetramethylheptane
and activate the hydrogen – hydrogen bond.
(ii) 2,5-Dimethyhexane Platinum and palladium catalyse the reaction
at room temperature but relatively higher
Solution
temperature and pressure are required with
nickel catalysts.
(i) CH3 – CH2 – CH – C – CH– CH – CH3 CH2 = CH2 + H2 ⎯⎯⎯⎯→
Pt/Pd/Ni
CH3 − CH3
(13.1)
Ethene Ethane
Propene Propane
(ii) CH3 – CH – CH2 – CH2 – CH – CH3
(13.2)
Problem 13.5
CH3 − C ≡ C − H + 2H2 ⎯⎯⎯⎯ → CH3 − CH2 − CH3
Pt/Pd/Ni
Write structures for each of the following
compounds. Why are the given names Propyne Propane
incorrect? Write correct IUPAC (13.3)
names.
2. From alkyl halides
(i) 2-Ethylpentane
i) Alkyl halides (except fluorides) on
(ii) 5-Ethyl – 3-methylheptane reduction with zinc and dilute hydrochloric
Solution acid give alkanes.
(i) CH3 – CH – CH2– CH2 – CH3 +
CH 3 − Cl + H 2 ⎯⎯ ⎯
Zn, H
→ CH 4 + HCl (13.4)
Chloromethane Methane
HYDROCARBONS 371
+
C 2 H 5 − Cl + H 2 ⎯⎯ ⎯
Zn, H
→ C 2 H 6 + HCl containing even number of carbon atoms
Chloroethane Ethane (13.5) at the anode.
+ 2CH3 COO− Na+ + 2H2 O
CH3 CH2 CH2 Cl + H2 ⎯⎯→
Zn,H
CH3 CH2 CH3 + HCl Sodium acetate
1-Chloropropane Propane
(13.6) ↓ Electrolysis
CH3 − CH3 + 2CO2 + H2 + 2NaOH (13.9)
ii) Alkyl halides on treatment with sodium
metal in dry ethereal (free from moisture) The reaction is supposed to follow the
solution give higher alkanes. This reaction following path :
is known as Wurtz reaction and is used O
for the preparation of higher alkanes ||
−
containing even number of carbon i) 2CH3 COO Na+ U 2CH3 −C −O− + 2Na+
atoms.
dry ether At anode:
CH3 Br +2Na + BrCH3 ⎯⎯⎯⎯ →CH3 −CH3 +2NaBr
O O
Bromomethane Ethane || ||
− • •
(13.7) –2e
2CH3 −C −O ⎯⎯
–
→2CH3 −C −O: ⎯⎯
→2CH3 +2CO2 ↑
••
dry ether
C2 H5 Br + 2Na + BrC2 H5 ⎯⎯⎯⎯ →C2 H5 − C2 H5
Acetate ion Acetate Methyl free
Bromoethane n-Butane
free radical radical
(13.8) • •
iii) H3 C + CH3 ⎯⎯
→ H3 C − CH3 ↑
What will happen if two different alkyl halides
are taken? iv) At cathode :
•
3. From carboxylic acids H2 O + e– → – OH + H
•
i) Sodium salts of carboxylic acids on heating 2H ⎯
→ H2 ↑
with soda lime (mixture of sodium
Methane cannot be prepared by this
hydroxide and calcium oxide) give alkanes
method. Why?
containing one carbon atom less than the
carboxylic acid. This process of elimination 13.2.3 Properties
of carbon dioxide from a carboxylic acid is Physical properties
known as decarboxylation. Alkanes are almost non-polar molecules
because of the covalent nature of C-C and C-H
CH3 COO Na+ + NaOH ⎯
–
⎯→CH4 + Na2 CO3
CaO
Δ bonds and due to very little difference of
Sodium ethanoate electronegativity between carbon and
hydrogen atoms. They possess weak van der
Problem 13.6 Waals forces. Due to the weak forces, the first
Sodium salt of which acid will be needed four members, C1 to C4 are gases, C5 to C17 are
for the preparation of propane ? Write liquids and those containing 18 carbon atoms
chemical equation for the reaction. or more are solids at 298 K. They are colourless
and odourless. What do you think about
Solution solubility of alkanes in water based upon non-
Butanoic acid, polar nature of alkanes? Petrol is a mixture of
−
CH3 CH2 CH2 COO Na + + NaOH ⎯⎯
CaO
→ hydrocarbons and is used as a fuel for
automobiles. Petrol and lower fractions of
CH3 CH2 CH3 + Na2 CO3
petroleum are also used for dry cleaning of
clothes to remove grease stains. On the basis
ii) Kolbe’s electrolytic method An aqueous of this observation, what do you think about
solution of sodium or potassium salt of a the nature of the greasy substance? You are
carboxylic acid on electrolysis gives alkane correct if you say that grease (mixture of higher
372 CHEMISTRY
alkanes) is non-polar and, hence, hydrophobic reducing agents. However, they undergo the
in nature. It is generally observed that in following reactions under certain
relation to solubility of substances in solvents, conditions.
polar substances are soluble in polar solvents,
1. Substitution reactions
whereas the non-polar ones in non-polar
solvents i.e., like dissolves like. One or more hydrogen atoms of alkanes can
be replaced by halogens, nitro group and
Boiling point (b.p.) of different alkanes are
sulphonic acid group. Halogenation takes
given in Table 13.2 from which it is clear that
there is a steady increase in boiling point with place either at higher temperature
increase in molecular mass. This is due to the (573-773 K) or in the presence of diffused
fact that the intermolecular van der Waals sunlight or ultraviolet light. Lower alkanes do
forces increase with increase of the molecular not undergo nitration and sulphonation
size or the surface area of the molecule. reactions. These reactions in which hydrogen
atoms of alkanes are substituted are known
You can make an interesting observation
as substitution reactions. As an example,
by having a look on the boiling points of
chlorination of methane is given below:
three isomeric pentanes viz., (pentane,
2-methylbutane and 2,2-dimethylpropane). It Halogenation
is observed (Table 13.2) that pentane having a hν
CH4 + Cl2 ⎯⎯ → CH3 Cl + HCl
continuous chain of five carbon atoms has the
highest boiling point (309.1K) whereas Chloromethane (13.10)
2,2 – dimethylpropane boils at 282.5K. With
hν
increase in number of branched chains, the CH 3 Cl + Cl2 ⎯⎯⎯
→ CH 2 Cl2 + HCl
molecule attains the shape of a sphere. This Dichloromethane (13.11)
results in smaller area of contact and therefore
weak intermolecular forces between spherical hν
CH2 Cl2 + Cl2 ⎯⎯⎯
→ CHCl3 + HCl
molecules, which are overcome at relatively Trichloromethane (13.12)
lower temperatures.
hν
Chemical properties CHCl3 + Cl2 ⎯⎯⎯
→ CCl4 + HCl
As already mentioned, alkanes are generally Tetrachloromethane (13.13)
inert towards acids, bases, oxidising and
Table 13.2 Variation of Melting Point and Boiling Point in Alkanes
13.3 ALKENES
Alkenes are unsaturated hydrocarbons
containing at least one double bond. What
should be the general formula of alkenes? If Fig. 13.4 Orbital picture of ethene depicting
there is one double bond between two carbon σ bonds only
atoms in alkenes, they must possess two
hydrogen atoms less than alkanes. Hence, 13.3.2 Nomenclature
general formula for alkenes is CnH2n. Alkenes For nomenclature of alkenes in IUPAC system,
are also known as olefins (oil forming) since the longest chain of carbon atoms containing
the first member, ethylene or ethene (C2H4) was the double bond is selected. Numbering of the
found to form an oily liquid on reaction with chain is done from the end which is nearer to
chlorine. the double bond. The suffix ‘ene’ replaces ‘ane’
HYDROCARBONS 377
Fig. 13.5 Orbital picture of ethene showing formation of (a) π-bond, (b) π-cloud and (c) bond angles
and bond lengths
But-1-ene
(iii) CH2 = C (CH2CH2CH3)2 (C4H8)
(iv) CH3 CH2 CH2 CH2 CH2CH3
| | II. 1 2 3 4
CH3 – CHCH = C – CH2 – CHCH3 CH3 – CH = CH – CH3
|
CH3 But-2-ene
(C4H8)
378 CHEMISTRY
13.3.4 Preparation
1. From alkynes: Alkynes on partial
reduction with calculated amount of
cis-But-2-ene trans-But-2-ene dihydrogen in the presence of palladised
(μ = 0.33D) (μ = 0) charcoal partially deactivated with poisons
In the case of solids, it is observed that like sulphur compounds or quinoline give
the trans isomer has higher melting point alkenes. Partially deactivated palladised
than the cis form. charcoal is known as Lindlar’s catalyst.
Alkenes thus obtained are having cis
Geometrical or cis-trans isomerism
geometry. However, alkynes on reduction
is also shown by alkenes of the types
XYC = CXZ and XYC = CZW with sodium in liquid ammonia form trans
alkenes.
Problem 13.10
Draw cis and trans isomers of the
following compounds. Also write their
IUPAC names :
(i) CHCl = CHCl (13.30)
(ii) C2H5CCH3 = CCH3C2H5
Solution
(13.31)
Propyne Propene
(13.33)
Will propene thus obtained show
geometrical isomerism? Think for the
Problem 13.11 reason in support of your answer.
Which of the following compounds will
show cis-trans isomerism? 2. From alkyl halides: Alkyl halides (R-X)
on heating with alcoholic potash
(i) (CH3)2C = CH – C2H5
(potassium hydroxide dissolved in alcohol,
380 CHEMISTRY
say, ethanol) eliminate one molecule of takes out one hydrogen atom from the
halogen acid to form alkenes. This reaction β-carbon atom.
is known as dehydrohalogenation i.e.,
removal of halogen acid. This is example of
β-elimination reaction, since hydrogen
atom is eliminated from the β carbon atom
(carbon atom next to the carbon to which
halogen is attached).
(13.37)
13.3.5 Properties
Physical properties
Alkenes as a class resemble alkanes in physical
properties, except in types of isomerism and
(13.34) difference in polar nature. The first three
members are gases, the next fourteen are
Nature of halogen atom and the alkyl
liquids and the higher ones are solids. Ethene
group determine rate of the reaction. It is
is a colourless gas with a faint sweet smell. All
observed that for halogens, the rate is:
other alkenes are colourless and odourless,
iodine > bromine > chlorine, while for alkyl
insoluble in water but fairly soluble in non-
groups it is : tert > secondary > primary.
polar solvents like benzene, petroleum ether.
3. From vicinal dihalides: Dihalides in They show a regular increase in boiling point
which two halogen atoms are attached to with increase in size i.e., every – CH2 group
two adjacent carbon atoms are known as added increases boiling point by 20–30 K. Like
vicinal dihalides. Vicinal dihalides on alkanes, straight chain alkenes have higher
treatment with zinc metal lose a molecule boiling point than isomeric branched chain
of ZnX2 to form an alkene. This reaction is compounds.
known as dehalogenation.
Chemical properties
CH2 Br − CH2 Br + Zn ⎯⎯
→CH2 = CH2 + ZnBr2 Alkenes are the rich source of loosely held
pi (π) electrons, due to which they show
(13.35)
addition reactions in which the electrophiles
CH3 CHBr − CH2 Br + Zn ⎯⎯
→ CH3CH = CH2 add on to the carbon-carbon double bond to
+ ZnBr2 form the addition products. Some reagents
also add by free radical mechanism. There are
(13.36)
cases when under special conditions, alkenes
4. From alcohols by acidic dehydration: also undergo free radical substitution
You have read during nomenclature of reactions. Oxidation and ozonolysis reactions
different homologous series in Unit 12 that are also quite prominent in alkenes. A brief
alcohols are the hydroxy derivatives of description of different reactions of alkenes is
alkanes. They are represented by R–OH given below:
where, R is CnH2n+1. Alcohols on heating
1. Addition of dihydrogen: Alkenes add up
with concentrated sulphuric acid form
one molecule of dihydrogen gas in the
alkenes with the elimination of one water
presence of finely divided nickel, palladium
molecule. Since a water molecule is
or platinum to form alkanes (Section 13.2.2)
eliminated from the alcohol molecule in the
presence of an acid, this reaction is known 2. Addition of halogens : Halogens like
as acidic dehydration of alcohols. This bromine or chlorine add up to alkene to
reaction is also the example of form vicinal dihalides. However, iodine
β-elimination reaction since –OH group does not show addition reaction under
HYDROCARBONS 381
Solution
• •
Homolysis
(ii) C6 H5 + H – Br ⎯⎯⎯⎯ ⎯→C6 H6 + Br
(13.49)
(13.44) KMnO /H+
CH3 – CH = CH – CH3 ⎯⎯⎯⎯⎯
4
→2CH3 COOH
But-2-ene Ethanoic acid
(13.50)
7. Ozonolysis : Ozonolysis of alkenes involves
the addition of ozone molecule to alkene to
form ozonide, and then cleavage of the
ozonide by Zn-H2O to smaller molecules.
This reaction is highly useful in detecting
(13.45) the position of the double bond in alkenes
or other unsaturated compounds.
5. Addition of water : In the presence of a
few drops of concentrated sulphuric acid
alkenes react with water to form alcohols,
in accordance with the Markovnikov rule.
(13.51)
(13.46)
6. Oxidation: Alkenes on reaction with cold,
dilute, aqueous solution of potassium
permanganate (Baeyer’s reagent) produce
vicinal glycols. Decolorisation of KMnO4
solution is used as a test for unsaturation.
(13.52)
(13.47) 8. Polymerisation: You are familiar with
polythene bags and polythene sheets.
Polythene is obtained by the combination
of large number of ethene molecules at high
temperature, high pressure and in the
presence of a catalyst. The large molecules
(13.48) thus obtained are called polymers. This
b) Acidic potassium permanganate or acidic reaction is known as polymerisation. The
potassium dichromate oxidises alkenes to simple compounds from which polymers
384 CHEMISTRY
are made are called monomers. Other are named as derivatives of the corresponding
alkenes also undergo polymerisation. alkanes replacing ‘ane’ by the suffix ‘yne’. The
High temp./pressure
position of the triple bond is indicated by the
n(CH2 = CH2 ) ⎯⎯⎯⎯⎯⎯⎯⎯
Catalyst → —( CH2 –CH2 —
)n first triply bonded carbon. Common and
Polythene IUPAC names of a few members of alkyne series
(13.53) are given in Table 13.2.
High temp./pressure
You have already learnt that ethyne and
n(CH3 – CH = CH2 ) ⎯⎯⎯⎯⎯⎯⎯⎯
Catalyst → —
( CH–CH2 —)n propyne have got only one structure but there
|
are two possible structures for butyne –
CH3
Polypropene (i) but-1-yne and (ii) but-2-yne. Since these two
compounds differ in their structures due to the
(13.54) position of the triple bond, they are known as
Polymers are used for the manufacture of position isomers. In how many ways, you can
plastic bags, squeeze bottles, refrigerator dishes, construct the structure for the next homologue
toys, pipes, radio and T.V. cabinets etc. i.e., the next alkyne with molecular formula
Polypropene is used for the manufacture of milk C5H8? Let us try to arrange five carbon atoms
crates, plastic buckets and other moulded with a continuous chain and with a side chain.
articles. Though these materials have now Following are the possible structures :
become common, excessive use of polythene
Structure IUPAC name
and polypropylene is a matter of great concern
for all of us. 1 2 3 4 5
I. HC ≡ C– CH – CH – CH Pent–1-yne
2 2 3
13.4 ALKYNES 1 2 3 4 5
II. H C– C ≡ C– CH – CH Pent–2-yne
3 2 3
Like alkenes, alkynes are also unsaturated
hydrocarbons. They contain at least one triple 4 3 2 1
III. H C– CH – C ≡ CH 3-Methyl but–1-yne
3
bond between two carbon atoms. The number
|
of hydrogen atoms is still less in alkynes as
CH3
compared to alkenes or alkanes. Their general
formula is CnH2n–2. Structures I and II are position isomers and
structures I and III or II and III are chain
The first stable member of alkyne series
isomers.
is ethyne which is popularly known as
acetylene. Acetylene is used for arc welding Problem 13.13
purposes in the form of oxyacetylene flame
Write structures of different isomers
obtained by mixing acetylene with oxygen gas. th
corresponding to the 5 member of
Alkynes are starting materials for a large
alkyne series. Also write IUPAC names of
number of organic compounds. Hence, it is
all the isomers. What type of isomerism is
interesting to study this class of organic
exhibited by different pairs of isomers?
compounds.
13.4.1 Nomenclature and Isomerism Solution
th
In common system, alkynes are named as 5 member of alkyne has the molecular
derivatives of acetylene. In IUPAC system, they formula C6H10. The possible isomers are:
Table 13.2 Common and IUPAC Names of Alkynes (CnH2n–2)
3-Methylpent-1-yne
4-Methylpent-1-yne
4-Methylpent-2-yne
(13.67)
(13.64)
Reddish orange colour of the solution of
bromine in carbon tetrachloride is decolourised.
This is used as a test for unsaturation.
(13.68)
(iii) Addition of hydrogen halides
(v) Polymerisation
Two molecules of hydrogen halides (HCl, HBr,
HI) add to alkynes to form gem dihalides (in (a) Linear polymerisation: Under suitable
which two halogens are attached to the same conditions, linear polymerisation of ethyne
carbon atom) takes place to produce polyacetylene or
H – C ≡ C – H + H – Br ⎯→ [CH2 = CH – Br] ⎯→ CHBr2 polyethyne which is a high molecular weight
Bromoethene | polyene containing repeating units of
CH3 (CH = CH – CH = CH ) and can be represented
1,1-Dibromoethane as —( CH = CH – CH = CH )— n Under special
(13.65) conditions, this polymer conducts electricity.
388 CHEMISTRY
(13.69)
Problem 13.14
How will you convert ethanoic acid into Biphenyl
benzene?
13.5.1 Nomenclature and Isomerism
Solution
The nomenclature and isomerism of aromatic
hydrocarbons has already been discussed in
Unit 12. All six hydrogen atoms in benzene are
equivalent; so it forms one and only one type
of monosubstituted product. When two
hydrogen atoms in benzene are replaced by
two similar or different monovalent atoms or
groups, three different position isomers are
possible. The 1, 2 or 1, 6 is known as the ortho
(o–), the 1, 3 or 1, 5 as meta (m–) and the 1, 4
as para (p–) disubstituted compounds. A few
examples of derivatives of benzene are given
below:
between all the carbon atoms in the ring has (i) Planarity
been determined by the X-ray diffraction to be (ii) Complete delocalisation of the π electrons
the same; there is equal probability for the p in the ring
orbital of each carbon atom to overlap with the
(iii) Presence of (4n + 2) π electrons in the ring
p orbitals of adjacent carbon atoms [Fig. 13.7
where n is an integer (n = 0, 1, 2, . . .).
(c)]. This can be represented in the form of two
doughtnuts (rings) of electron clouds [Fig. 13.7 This is often referred to as Hückel Rule.
(d)], one above and one below the plane of the Some examples of aromatic compounds are
hexagonal ring as shown below: given below:
(electron cloud)
(iii) Reduction of phenol: Phenol is reduced (ii) Halogenation: Arenes react with halogens
to benzene by passing its vapours over in the presence of a Lewis acid like anhydrous
heated zinc dust FeCl3, FeBr3 or AlCl3 to yield haloarenes.
(13.71) Chlorobenzene
13.5.5 Properties (13.73)
Physical properties (iii) Sulphonation: The replacement of a
Aromatic hydrocarbons are non- polar hydrogen atom by a sulphonic acid group in
molecules and are usually colourless liquids a ring is called sulphonation. It is carried out
or solids with a characteristic aroma. You are by heating benzene with fuming sulphuric acid
also familiar with naphthalene balls which are (oleum).
used in toilets and for preservation of clothes
because of unique smell of the compound and
the moth repellent property. Aromatic
hydrocarbons are immiscible with water but
are readily miscible with organic solvents. They
burn with sooty flame.
Chemical properties (13.74)
Arenes are characterised by electrophilic
substitution reactions. However, under special (iv) Friedel-Crafts alkylation reaction:
conditions they can also undergo addition and When benzene is treated with an alkyl halide
oxidation reactions. in the presence of anhydrous aluminium
chloride, alkylbenene is formed.
Electrophilic substitution reactions
The common electrophilic substitution
reactions of arenes are nitration, halogenation,
sulphonation, Friedel Craft’s alkylation and
acylation reactions in which attacking reagent
+
is an electrophile (E )
(i) Nitration: A nitro group is introduced into (13.75)
benzene ring when benzene is heated with a
mixture of concentrated nitric acid and
concentrated sulphuric acid (nitrating
mixture).
(13.76)
(13.77)
Benzene hexachloride,
(BHC)
(13.81) It is clear from the above resonating structures
Combustion: When heated in air, benzene that the electron density is more on
burns with sooty flame producing CO2 and o – and p – positions. Hence, the substitution
H2O takes place mainly at these positions. However,
it may be noted that –I effect of – OH group
15
C6 H 6 + O2 → 6CO2 + 3H2O (13.82) also operates due to which the electron density
2 on ortho and para positions of the benzene ring
General combustion reaction for any is slightly reduced. But the overall electron
hydrocarbon may be given by the following density increases at these positions of the ring
HYDROCARBONS 395
due to resonance. Therefore, –OH group are also called ‘deactivating groups’. The
activates the benzene ring for the attack by electron density on o – and p – position is
an electrophile. Other examples of activating comparatively less than that at meta position.
groups are –NH2, –NHR, –NHCOCH3, –OCH3, Hence, the electrophile attacks on
–CH3, –C2H5, etc. comparatively electron rich meta position
In the case of aryl halides, halogens are resulting in meta substitution.
moderately deactivating. Because of their
13.6 CARCINOGENICITY AND TOXICITY
strong – I effect, overall electron density on
benzene ring decreases. It makes further Benzene and polynuclear hydrocarbons
substitution difficult. However, due to containing more than two benzene rings
resonance the electron density on o – and p– fused together are toxic and said to possess
positions is greater than that at the m-position. cancer producing (carcinogenic) property.
Hence, they are also o – and p – directing Such polynuclear hydrocarbons are formed
groups. on incomplete combustion of organic
materials like tobacco, coal and petroleum.
Meta directing group: The groups which
They enter into human body and undergo
direct the incoming group to meta position are
various biochemical reactions and finally
called meta directing groups. Some examples
damage DNA and cause cancer. Some of
of meta directing groups are –NO2, –CN, –CHO,
the carcinogenic hydrocarbons are given
–COR, –COOH, –COOR, –SO3H, etc.
below (see box).
Let us take the example of nitro group. Nitro
group reduces the electron density in the
benzene ring due to its strong –I effect.
Nitrobenzene is a resonance hybrid of the
following structures.
SUMMARY
Hydrocarbons are the compounds of carbon and hydrogen only. Hydrocarbons are mainly
obtained from coal and petroleum, which are the major sources of energy.
Petrochemicals are the prominent starting materials used for the manufacture of a
large number of commercially important products. LPG (liquefied petroleum gas) and
CNG (compressed natural gas), the main sources of energy for domestic fuels and the
automobile industry, are obtained from petroleum. Hydrocarbons are classified as open
chain saturated (alkanes) and unsaturated (alkenes and alkynes), cyclic (alicyclic)
and aromatic, according to their structure.
The important reactions of alkanes are free radical substitution, combustion,
oxidation and aromatization. Alkenes and alkynes undergo addition reactions, which
are mainly electrophilic additions. Aromatic hydrocarbons, despite having unsaturation,
undergo mainly electrophilic substitution reactions. These undergo addition reactions
only under special conditions.
Alkanes show conformational isomerism due to free rotation along the C–C sigma
bonds. Out of staggered and the eclipsed conformations of ethane, staggered conformation
is more stable as hydrogen atoms are farthest apart. Alkenes exhibit geometrical
(cis-trans) isomerism due to restricted rotation around the carbon–carbon double bond.
Benzene and benzenoid compounds show aromatic character. Aromaticity, the
property of being aromatic is possessed by compounds having specific electronic structure
characterised by Hückel (4n+2)π electron rule. The nature of groups or substituents
attached to benzene ring is responsible for activation or deactivation of the benzene ring
towards further electrophilic substitution and also for orientation of the incoming group.
Some of the polynuclear hydrocarbons having fused benzene ring system have
carcinogenic property.
EXERCISES
13.1 How do you account for the formation of ethane during chlorination of methane ?
13.2 Write IUPAC names of the following compounds :
(a) CH3CH=C(CH3)2 (b) CH2=CH-C≡C-CH3
UNIT 14
ENVIRONMENTAL CHEMISTRY
(c) Hydrocarbons: Hydrocarbons are atmosphere. With the increased use of fossil
composed of hydrogen and carbon only and fuels, a large amount of carbon dioxide gets
are formed by incomplete combustion of fuel released into the atmosphere. Excess of CO2
used in automobiles. Hydrocarbons are in the air is removed by green plants and this
carcinogenic, i.e., they cause cancer. They maintains an appropriate level of CO2 in the
harm plants by causing ageing, breakdown of atmosphere. Green plants require CO2 for
tissues and shedding of leaves, flowers and photosynthesis and they, in turn, emit oxygen,
twigs. thus maintaining the delicate balance. As you
(d) Oxides of Carbon know, deforestation and burning of fossil fuel
(i ) Carbon monoxide: Carbon monoxide (CO) increases the CO2 level and disturb the balance
is one of the most serious air pollutants. It is a in the atmosphere. The increased amount of
colourless and odourless gas, highly CO2 in the air is mainly responsible for global
poisonous to living beings because of its ability warming.
to block the delivery of oxygen to the organs Global Warming and Greenhouse Effect
and tissues. It is produced as a result of About 75 % of the solar energy reaching the
incomplete combustion of carbon. Carbon earth is absorbed by the earth’s surface, which
monoxide is mainly released into the air by increases its temperature. The rest of the heat
automobile exhaust. Other sources, which radiates back to the atmosphere. Some of the
produce CO, involve incomplete combustion
heat is trapped by gases such as carbon
of coal, firewood, petrol, etc. The number of
dioxide, methane, ozone, chlorofluorocarbon
vehicles has been increasing over the years all
compounds (CFCs) and water vapour in the
over the world. Many vehicles are poorly
atmosphere. Thus, they add to the heating of
maintained and several have inadequate
the atmosphere. This causes global warming.
pollution control equipments resulting in the
release of greater amount of carbon monoxide We all know that in cold places flowers,
and other polluting gases. Do you know why vegetables and fruits are grown in glass
carbon monoxide is poisonous? It binds to covered areas called greenhouse. Do you know
haemoglobin to form carboxyhaemoglobin, that we humans also live in a greenhouse? Of
which is about 300 times more stable than the course, we are not surrounded by glass but a
oxygen-haemoglobin complex. In blood, when blanket of air called the atmosphere, which has
the concentration of carboxyhaemoglobin kept the temperature on earth constant for
reaches about 3–4 per cent, the oxygen centuries. But it is now undergoing change,
carrying capacity of blood is greatly though slowly. Just as the glass in a
reduced. This oxygen deficiency, results into greenhouse holds the sun’s warmth inside,
headache, weak eyesight, nervousness and atmosphere traps the sun’s heat near the
cardiovascular disorder. This is the reason why earth’s surface and keeps it warm. This is
people are advised not to smoke. In pregnant called natural greenhouse effect because it
women who have the habit of smoking the maintains the temperature and makes the
increased CO level in blood may induce earth perfect for life. In a greenhouse, visible
premature birth, spontaneous abortions and light passes through the transparent glass and
deformed babies. heats up the soil and the plants. The warm
(ii) Carbon dioxide: Carbon dioxide (CO2) is soil and plants emit infrared radiations. Since
released into the atmosphere by respiration, glass is opaque to infrared (heat) radiations, it
burning of fossil fuels for energy, and by partly reflects and partly absorbs these
decomposition of limestone during the radiations. This mechanism keeps the energy
manufacture of cement. It is also emitted of the sun trapped in the greenhouse.
during volcanic eruptions. Carbon dioxide gas Similarly, carbon dioxide molecules also trap
is confined to troposphere only. Normally it heat as they are transparent to sunlight but
forms about 0.03 per cent by volume of the not to the heat radiation. If the amount of
ENVIRONMENTAL CHEMISTRY 401
This plan aims at clearing the air in herbicides and insecticides that miss their
the ‘Taj Trapezium’– an area that includes targets and travel through air and form
the towns of Agra, Firozabad, Mathura and mists.
Bharatpur. Under this plan more than
(d) Fumes are generally obtained by the
2000 polluting industries lying inside the
trapezium would switch over to the use of condensation of vapours during
natural gas or liquefied petroleum gas sublimation, distillation, boiling and
instead of coal or oil. A new natural gas several other chemical reactions. Generally,
pipeline would bring more than half a organic solvents, metals and metallic
million cubic metres of natural gas a day oxides form fume particles.
to this area. People living in the city will
also be encouraged to use liquefied
The effect of particulate pollutants are
petroleum gas in place of coal, kerosene or largely dependent on the particle size. Air-
firewood. Vehicles plying on highways in borne particles such as dust, fumes, mist etc.,
the vicinity of Taj would be encouraged to are dangerous for human health. Particulate
use low sulphur content diesel. pollutants bigger than 5 microns are likely to
lodge in the nasal passage, whereas particles
2. Particulate Pollutants of about 1.0 micron enter into lungs easily.
Particulates pollutants are the minute solid Lead used to be a major air pollutant
particles or liquid droplets in air. These are emitted by vehicles. Leaded petrol used to be
present in vehicle emissions, smoke particles the primary source of air-borne lead emission
from fires, dust particles and ash from in Indian cities. This problem has now been
industries. Particulates in the atmosphere overcome by using unleaded petrol in most of
may be viable or non-viable. The viable the cities in India. Lead interferes with the
particulates e.g., bacteria, fungi, moulds, development and maturation of red blood cells.
algae etc., are minute living organisms that are Smog
dispersed in the atmosphere. Human beings The word smog is derived from smoke and fog.
are allergic to some of the fungi found in air. This is the most common example of air
They can also cause plant diseases. pollution that occurs in many cities
Non-viable particulates may be classified throughout the world. There are two types of
according to their nature and size as follows: smog:
(a) Smoke particulates consist of solid or (a) Classical smog occurs in cool humid
mixture of solid and liquid particles formed climate. It is a mixture of smoke, fog and
during combustion of organic matter. sulphur dioxide. Chemically it is a
Examples are cigarette smoke, smoke from reducing mixture and so it is also called
burning of fossil fuel, garbage and dry as reducing smog.
leaves, oil smoke etc. (b) Photochemical smog occurs in warm, dry
(b) Dust is composed of fine solid particles and sunny climate. The main components
(over 1µm in diameter), produced during of the photochemical smog result from the
crushing, grinding and attribution of solid action of sunlight on unsaturated
materials. Sand from sand blasting, saw hydrocarbons and nitrogen oxides
dust from wood works, pulverized coal, produced by automobiles and factories.
cement and fly ash from factories, dust Photochemical smog has high
storms etc., are some typical examples of concentration of oxidising agents and is,
this type of particulate emission. therefore, called as oxidising smog.
(c) Mists are produced by particles of spray Formation of photochemical smog
liquids and by condensation of vapours in When fossil fuels are burnt, a variety of
air. Examples are sulphuric acid mist and pollutants are emitted into the earth’s
404 CHEMISTRY
troposphere. Two of the pollutants that are to produce chemicals such as formaldehyde,
emitted are hydrocarbons (unburnt fuels) and acrolein and peroxyacetyl nitrate (PAN).
nitric oxide (NO). When these pollutants build 3CH4 + 2O3 → 3CH2 = O + 3H2O
up to sufficiently high levels, a chain reaction
Formaldehyde
occurs from their interaction with sunlight in
which NO is converted into nitrogen dioxide CH2=CHCH=O CH3COONO2
(NO2). This NO2 in turn absorbs energy from
sunlight and breaks up into nitric oxide and O
free oxygen atom (Fig. 14.2). Acrolein Peroxyacetyl nitrate (PAN)
NO2(g) NO(g) + O(g) (i) Effects of photochemical smog
Oxygen atoms are very reactive and The common components of photochemical
combine with the O2 in air to produce ozone. smog are ozone, nitric oxide, acrolein,
formaldehyde and peroxyacetyl nitrate (PAN).
O(g) + O2 (g) U O3 (g) (ii) Photochemical smog causes serious health
The ozone formed in the above reaction (ii) problems. Both ozone and PAN act as powerful
reacts rapidly with the NO(g) formed in the eye irritants. Ozone and nitric oxide irritate the
reaction (i) to regenerate NO2. NO2 is a brown nose and throat and their high concentration
gas and at sufficiently high levels can causes headache, chest pain, dryness of the
contribute to haze. throat, cough and difficulty in breathing.
Photochemical smog leads to cracking of
NO (g) + O3 (g) → NO2 (g) + O2 (g) (iii) rubber and extensive damage to plant life. It
Ozone is a toxic gas and both NO2 and O3 also causes corrosion of metals, stones,
are strong oxidising agents and can react with building materials, rubber and painted
the unburnt hydrocarbons in the polluted air surfaces.
Fig. 14.2 Photochemical smog occurs where sunlight acts on vehicle pollutants.
ENVIRONMENTAL CHEMISTRY 405
How can photochemical smog be in the production of plastic foam and by the
controlled ? electronic industry for cleaning computer
Many techniques are used to control or reduce parts etc. Once CFCs are released in the
the formation of photochemical smog. If we atmosphere, they mix with the normal
control the primary precursors of atmospheric gases and eventually reach the
photochemical smog, such as NO 2 and stratosphere. In stratosphere, they get broken
hydrocarbons, the secondary precursors such down by powerful UV radiations, releasing
as ozone and PAN, the photochemical smog chlorine free radical.
will automatically be reduced. Usually catalytic
CF2Cl2 (g) (g) + F2Cl (g) (i)
converters are used in the automobiles, which
prevent the release of nitrogen oxide and The chlorine radical then react with
hydrocarbons to the atmosphere. Certain stratospheric ozone to form chlorine monoxide
plants e.g., Pinus, Juniparus, Quercus, Pyrus radicals and molecular oxygen.
• •
and Vitis can metabolise nitrogen oxide and C l (g) + O3 (g) → Cl O (g) + O2 (g) (ii)
therefore, their plantation could help in this
matter. Reaction of chlorine monoxide radical with
atomic oxygen produces more chlorine
14.2.2 Stratospheric Pollution radicals.
• •
Formation and Breakdown of Ozone Cl O (g) + O (g) → C l (g) + O2 (g) (iii)
The upper stratosphere consists of The chlorine radicals are continuously
considerable amount of ozone (O3), which regenerated and cause the breakdown of
protects us from the harmful ultraviolet (UV)
ozone. Thus, CFCs are transporting agents for
radiations (λ 255 nm) coming from the sun.
continuously generating chlorine radicals into
These radiations cause skin cancer
the stratosphere and damaging the ozone layer.
(melanoma) in humans. Therefore, it is
important to maintain the ozone shield. The Ozone Hole
Ozone in the stratosphere is a product of In 1980s atmospheric scientists working in
UV radiations acting on dioxygen (O 2 ) Antarctica reported about depletion of ozone
molecules. The UV radiations split apart layer commonly known as ozone hole over the
molecular oxygen into free oxygen (O) atoms. South Pole. It was found that a unique set of
These oxygen atoms combine with the conditions was responsible for the ozone hole.
molecular oxygen to form ozone. In summer season, nitrogen dioxide and
methane react with chlorine monoxide
O2 (g) O(g) + O(g) (reaction iv) and chlorine atoms (reaction v)
O(g) + O2 (g) O3 (g) forming chlorine sinks, preventing much ozone
depletion, whereas in winter, special type of
Ozone is thermodynamically unstable and
clouds called polar stratospheric clouds are
decomposes to molecular oxygen. Thus, a
formed over Antarctica. These polar
dynamic equilibrium exists between the
stratospheric clouds provide surface on which
production and decomposition of ozone
chlorine nitrate formed (reaction iv) gets
molecules. In recent years, there have been
hydrolysed to form hypochlorous acid
reports of the depletion of this protective ozone
(reaction (vi)). It also reacts with hydrogen
layer because of the presence of certain
chloride produced as per reaction (v) to give
chemicals in the stratosphere. The main
molecular chlorine.
reason of ozone layer depletion is believed to •
be the release of chlorofluorocarbon Cl O (g) + NO2 (g) → ClONO2(g) (iv)
• •
compounds (CFCs), also known as freons. C l (g) + CH4 (g) → C H3(g) + HCl(g) (v)
These compounds are nonreactive, non
ClONO2(g) + H2O (g) → HOCl (g) + HNO3 (g) (vi)
flammable, non toxic organic molecules and
therefore used in refrigerators, air conditioners, ClONO2(g) + HCl (g) → Cl2 (g) + HNO3 (g) (vii)
406 CHEMISTRY
When sunlight returns to the Antarctica in where pollutants enter the water-source. Non
the spring, the sun’s warmth breaks up the point sources of pollution are those where a
clouds and HOCl and Cl2 are photolysed by source of pollution cannot be easily identified,
sunlight, as given in reactions (viii) and (ix). e.g., agricultural run off (from farm, animals
• • and crop-lands), acid rain, storm-water
HOCl (g)
hν
→ O H (g) + Cl(g) (viii)
drainage (from streets, parking lots and lawns),
•
Cl2 (g)
hν
→ 2 Cl (g) (ix) etc. Table 14.1 lists the major water pollutants
and their sources.
The chlorine radicals thus formed, initiate
the chain reaction for ozone depletion as 14.3.1 Causes of Water Pollution
described earlier. (i) Pathogens: The most serious water
Effects of Depletion of the Ozone Layer pollutants are the disease causing agents
called pathogens. Pathogens include bacteria
With the depletion of ozone layer, more UV
and other organisms that enter water from
radiation filters into troposphere. UV
domestic sewage and animal excreta. Human
radiations lead to ageing of skin, cataract,
excreta contain bacteria such as Escherichia
sunbur n, skin cancer, killing of many
coli and Streptococcus faecalis which cause
phytoplanktons, damage to fish productivity
gastrointestinal diseases.
etc. It has also been reported that plant
proteins get easily affected by UV radiations (ii) Organic wastes: The other major water
which leads to the harmful mutation of cells. pollutant is organic matter such as
It also increases evaporation of surface water leaves, grass, trash etc. They pollute water as
through the stomata of the leaves and a consequence of run off. Excessive
decreases the moisture content of the soil. phytoplankton growth within water is also a
Increase in UV radiations damage paints and cause of water pollution. These wastes are
fibres, causing them to fade faster. biodegradable.
The large population of bacteria
14.3 WATER POLLUTION decomposes organic matter present in water.
Water is essential for life. Without water there They consume oxygen dissolved in water. The
would be no life. We usually take water as amount of oxygen that water can hold in the
granted for its purity, but we must ensure the solution is limited. In cold water, dissolved
quality of water. Pollution of water originates oxygen (DO) can reach a concentration up to
from human activities. Through different 10 ppm (parts per million), whereas oxygen in
paths, pollution reaches surface or ground air is about 200,000 ppm. That is why even a
water. Easily identified source or place of moderate amount of organic matter when
pollution is called as point source. e.g., decomposes in water can deplete the water of
municipal and industrial discharge pipes its dissolved oxygen. The concentration of
Pollutant Source
Micro-organisms Domestic sewage
Organic wastes Domestic sewage, animal excreta and waste, decaying animals
and plants, discharge from food processing factories.
Plant nutrients Chemcial fertilizers
Toxic heavy metals Industries and chemical factories
Sediments Erosion of soil by agriculture and strip mining
Pesticides Chemicals used for killing insects, fungi and weeds
Radioactive substances Mining of uranium containing minerals
Heat Water used for cooling in industries
ENVIRONMENTAL CHEMISTRY 407
dissolved oxygen in water is very important The organic chemicals are another group
for aquatic life . If the concentration of dissolved of substances that are found in polluted water.
oxygen of water is below 6 ppm, the growth of Petroleum products pollute many sources of
fish gets inhibited. Oxygen reaches water water e.g., major oil spills in oceans. Other
either through atmosphere or from the process organic substances with serious impacts are
of photosynthesis carried out by many the pesticides that drift down from sprays or
aquatic green plants during day light. runof f from lands. Various industrial
However, during night, photosynthesis stops chemicals like polychlorinated biphenyls,
but the plants continue to respire, resulting (PCBs) which are used as cleansing solvent,
in reduction of dissolved oxygen. The detergents and fertilizers add to the list of
dissolved oxygen is also used by water pollutants. PCBs are suspected to be
microorganisms to oxidise organic matter. carcinogenic. Nowadays most of the detergents
available are biodegradable. However, their use
If too much of organic matter is added to
can create other problems. The bacteria
water, all the available oxygen is used up. This
responsible for degrading biodegradable
causes oxygen dependent aquatic life to die.
detergent feed on it and grow rapidly. While
Thus, anaerobic bacteria (which do not require
growing, they may use up all the oxygen
oxygen) begin to break down the organic waste
dissolved in water. The lack of oxygen kills all
and produce chemicals that have a foul smell
other forms of aquatic life such as fish and
and are harmful to human health. Aerobic
plants. Fertilizers contain phosphates as
(oxygen requiring) bacteria degrade these
additives. The addition of phosphates in water
organic wastes and keep the water depleted
enhances algae growth. Such profuse growth
in dissolved oxygen. of algae, covers the water surface and reduces
Thus, the amount of oxygen required by the oxygen concentration in water. This leads
bacteria to break down the organic matter to anaerobic conditions, commonly with
present in a certain volume of a sample of accumulation of abnoxious decay and animal
water, is called Biochemical Oxygen Demand death. Thus, bloom-infested water inhibits the
(BOD). The amount of BOD in the water is a growth of other living organisms in the
measure of the amount of organic material in water body. This process in which nutrient
the water, in terms of how much oxygen will enriched water bodies support a dense plant
be required to break it down biologically. Clean population, which kills animal life by depriving
water would have BOD value of less than it of oxygen and results in subsequent loss of
5 ppm whereas highly polluted water could biodiversity is known as Eutrophication.
have a BOD value of 17 ppm or more.
14.3.2 International Standards for
(iii) Chemical Pollutants: As we know that Drinking Water
water is an excellent solvent, water soluble
The International Standards for drinking water
inorganic chemicals that include heavy metals
are given below and they must be followed.
such as cadmium, mercury, nickel etc
constitute an important class of pollutants. All Fluoride: For drinking purposes, water
these metals are dangerous to humans should be tested for fluoride ion concentration.
because our body cannot excrete them. Over Its deficiency in drinking water is harmful to
the time, it crosses the tolerance limit. These man and causes diseases such as tooth decay
metals then can damage kidneys, central etc. Soluble fluoride is often added to drinking
nervous system, liver etc. Acids (like sulphuric water to bring its concentration upto 1 ppm
–3 –
acid) from mine drainage and salts from many or 1 mg dm . The F ions make the enamel on
different sources including raw salt used to teeth much harder by converting
melt snow and ice in the colder climates hydroxyapatite, [3(Ca3(PO4)2.Ca(OH)2], the
(sodium and calcium chloride) are water enamel on the surface of the teeth, into much
soluble chemical pollutants. harder fluorapatite, [3(Ca 3 (PO 4 ) 2 .CaF 2 ].
408 CHEMISTRY
–
However, F ion concentration above 2 ppm pollution levels. Ensure that appropriate
causes brown mottling of teeth. At the same action is taken. You can write to the press
time, excess fluoride (over 10 ppm) causes also. Do not dump waste into a
harmful effect to bones and teeth, as reported household or industrial drain which can
from some parts of Rajasthan. enter directly to any water body, such as,
Lead: Drinking water gets contaminated with river, pond, stream or lake. Use compost
lead when lead pipes are used for instead of chemical fertilizers in gardens.
transportation of water. The prescribed upper Avoid the use of pesticides like DDT,
limit concentration of lead in drinking water malathion etc., at home and try to use
is about 50 ppb. Lead can damage kidney, dried neem leaves to help keep insects
liver, reproductive system etc. away. Add a few crystals of potassium
permanganate (KMnO 4) or bleaching
Sulphate: Excessive sulphate (>500 ppm) in powder to the water tank of your house.
drinking water causes laxative effect, otherwise
at moderate levels it is harmless.
14.4 SOIL POLLUTION
Nitrate: The maximum limit of nitrate in
India being an agriculture based economy
drinking water is 50 ppm. Excess nitrate in
gives high priority to agriculture, fisheries and
drinking water can cause disease such as
livestock development. The surplus
methemoglobinemia (‘blue baby’ syndrome).
production is stored by governmental and
Other metals: The maximum concentration non-governmental organisations for the lean
of some common metals recommended in season. The food loss during the storage also
drinking water are given in Table 14.2. needs special attention. Have you ever seen the
damages caused to the crops, food items by
insects, rodents, weeds and crop diseases etc?
Table 14.2 Maximum Prescribed Concen- How can we protect them? You are acquainted
tration of Some Metals in
with some insecticides and pesticides for
Drinking Water.
protection of our crops. However, these
Metal Maximum concentration insecticides, pesticides and herbicides cause
–3
(ppm or mg dm ) soil pollution. Hence, there is a need for their
Fe 0.2 judicious use.
Mn 0.05 14.4.1 Pesticides
Al 0.2 Prior to World War II, many naturally
Cu 3.0 occurring chemicals such as nicotine (by
Zn 5.0 planting tobacco plants in the crop field), were
Cd 0.005 used as pest controlling substance for major
crops in agricultural practices.
Activity 2 During World War II, DDT was found to be
of great use in the control of malaria and other
You can visit local water sources and insect-borne diseases. Therefore, after the war,
observe if the river/lake/tank/pond are DDT was put to use in agriculture to control
unpolluted/slightly polluted/ moderately the damages caused by insects, rodents, weeds
polluted or severely polluted by looking and various crop diseases. However, due to
at water or by checking pH of water. adverse effects, its use has been banned in
Document the name of the river and the India.
nearby urban or industrial site from
where the pollution is generated. Inform Pesticides are basically synthetic toxic
about this to Pollution Control Board’s chemicals with ecological repercussions. The
office set up by Government to measure repeated use of the same or similar pesticides
give rise to pests that are resistant to that
ENVIRONMENTAL CHEMISTRY 409
group of pesticides thus making the pesticides sodium chlorate (NaClO3), sodium arsinite
ineffective. Therefore, as insect resistance of (Na3AsO3) and many others. During the first
DDT increased, other organic toxins such as half of the last century, the shift from
Aldrin and Dieldrin were introduced in the mechanical to chemical weed control had
market by pesticide industry. Most of the provided the industry with flourishing
organic toxins are water insoluble and non- economic market. But one must remember that
biodegradable. These high persistent toxins these are also not environment friendly.
are, therefore, transferred from lower trophic Most herbicides are toxic to mammals but
level to higher trophic level through food chain are not as persistent as organo-chlorides.
(Fig.14.3). Over the time, the concentration of These chemicals decompose in a few months.
toxins in higher animals reach a level which Like organo-chlorides, these too become
causes serious metabolic and physiological concentrated in the food web. Some herbicides
disorders. cause birth defects. Studies show that corn-
fields sprayed with herbicides are more prone
to insect attack and plant disease than fields
that are weeded manually.
Pesticides and herbicides represent only a
very small portion of widespread chemical
pollution. A large number of other compounds
that are used regularly in chemical and
industrial processes for manufacturing
activities are finally released in the atmosphere
in one or other form.
quantities of toxic wastes are usually destroyed household discards, there are medical,
by controlled incineration, whereas small agricultural, industrial and mining wastes. The
quantities are burnt along with factory improper disposal of wastes is one of the major
garbage in open bins. Moreover, solid wastes causes of environmental degradation.
if not managed effectively, affect the Therefore, the management of wastes is of
components of the environment. utmost importance.
Collection and Disposal
Do you know about waste recycling?
Domestic wastes are collected in small bins,
• Fuel obtained from plastic waste has which are then transferred to community bins
high octane rating. It contains no lead by private or municipal workers. From these
and is known as “green fuel”. community bins, these are collected and
• Due to recent developments made in carried to the disposable site. At the site,
chemical and textile industries, clothes garbage is sorted out and separated into
will be made from recycled plastic biodegradable and non-biodegradable
waste. These will be available soon in materials. Non-biodegradable materials such
the global textile market. as plastic, glass, metal scraps etc. are sent for
• In India, our cities and towns face recycling. Biodegradable wastes are deposited
endless hours of power cut. We can also in land fills and are converted into compost.
see piles of rotting garbage here and
The waste if not collected in garbage bins,
there. There is a good news that we can
finds its way into the sewers. Some of it is eaten
get rid from both these problems
by cattle. Non-biodegradable wastes like
simultaneously. Technology has now
been developed to produce electricity
polythene bag, metal scraps, etc. choke the
from the garbage. A pilot plant has been sewers and cause inconvenience. Polythene
set up, where after removing ferrous bags, if swallowed by cattle can cost their lives
metals, plastic, glass, paper etc. from also.
garbage, it is mixed with water. It is then As a normal practice, therefore, all
cultured with bacterial species for domestic wastes should be properly collected
producing methane, commonly known and disposed. The poor management causes
as biogas. The remaining product is health problems leading to epidemics due to
used as manure and biogas is used to contamination of ground water. It is specially
produce electricity. hazardous for those who are in direct contact
with the waste such as rag pickers and workers
14.6 STRATEGIES TO CONTROL involved in waste disposal, as they are the ones
who handle waste materials mostly without
ENVIRONMENTAL POLLUTION
protective device such as gloves or water proof
After studying air, water, soil and industrial boots and gas masks. What can you do for
waste pollution in this unit, by now you must them?
have started feeling the need of controlling
environmental pollution: How can you save 14.7 GREEN CHEMISTRY
your immediate environment? Think of the 14.7.1 Introduction
steps/activities, which you would like to
It is well known fact that self-sufficiency in food
undertake for controlling air, water, soil and th
industrial waste pollution in your has been achieved in India since late 20
neighbourhood. Here, an idea about the century by using fertilizers and pesticides and
strategies for the management of waste is given. exploring improved methods of farming, good
quality seeds, irrigation etc. But over -
14.6.1 Waste Management exploitation of soil and excessive use of
Solid waste is not the only waste, which you fertilizers and pesticides have resulted in the
see in your household garbage box. Besides deterioration of soil, water and air.
ENVIRONMENTAL CHEMISTRY 411
The solution of this problem does not lie in tetrachloride etc., are highly toxic. One should
stopping the process of development that has be careful while using them.
been set in; but to discover methods, which
As you know, a chemical reaction involves
would help in the reduction of deterioration of
reactants, attacking reagents and the medium
the environment. Green chemistry is a way of
in which the reaction takes place. Extent of any
thinking and is about utilising the existing
knowledge and principles of chemistry and reaction depends upon physical parameters
other sciences to reduce the adverse impact like temperature, pressure and use of catalyst.
on environment. Green chemistry is a In a chemical reaction, if reactants are fully
production process that would bring about converted into useful environmental friendly
minimum pollution or deterioration to the products by using an environment friendly
environment. The byproducts generated medium then there would be no chemical
during a process, if not used gainfully, add pollutants introduced in the environment.
to the environmental pollution. Such During a synthesis, care must be taken to
processes are not only environmental
choose starting materials that can be converted
unfriendly but also cost-ineffective. The
into end products with yield approximately
waste generation and its disposal both are
economically unsound. Utilisation of existing upto 100 per cent. This can be achieved by
knowledge base for reducing the chemical arriving at optimum conditions of synthesis.
hazards along with the developmental It may be worthwhile to carry out synthetic
activities is the foundation of green reactions in aqueous medium since water has
chemistry. Have you perceived the idea of green high specific heat and low volatility. Water is
chemistry ? It is well known that organic cost effective, noninflammable and devoid of
solvents such as benzene, toluene, carbon any carcinogenic effects.
SUMMARY
energy for the sustainance of life. The increase in the greenhouse gases is raising the
temperature of the earth’s atmosphere which, if not checked, may eventually result in
melting of polar ice caps and consequently may submerge the costal land mass. Many
human activities are producing chemicals, which are responsible for the depletion of
ozone layer in the stratosphere, leading to the formation of ozone hole. Through the
ozone hole, ultraviolet radiations can penetrate into the earth’s atmosphere causing
mutation of genes. Water is the elixir of life but the same water, if polluted by pathogens,
organic wastes, toxic heavy metals, pesticides etc., will turn into poison. Therefore, one
should take care to follow international standards to maintain purity levels of drinking
water. Industrial wastes and excessive use of pesticides, result into pollution of land
mass and water bodies. Judicious use of chemicals required for agricultural practices
can lead to sustainable development. Strategies for controlling environmental
pollution can be: (i) waste management i.e., reduction of the waste and proper disposal,
also recycling of materials and energy, (ii) adopting methods in day-to-day life, which
results in the reduction of environmental pollution. The second method is a new branch
of chemistry, which is in its infancy known as green chemistry. It utilizes the existing
knowledge and practices so as to bring about reduction in the production of pollutants.
EXERCISES