Basic Concepts P-Block Class 12
Basic Concepts P-Block Class 12
Basic Concepts P-Block Class 12
Group 15 Element-Introduction
Group 15 of the long form periodic table consists of five elements: Nitrogen (N), Phosphorus (P),
Arsenic (As), Antimony (Sb) and Bismuth (Bi). These elements are collectively called the elements of
nitrogen family. (Nicogen family)
Nitrogen and phosphorus are non-metals; arsenic and antimony are semi-metals or metalloids,
whereas bismuth is metallic although not so strongly.
bottom of the group. Thus, for heavier elements of the group (e.g., Sb and Bi.), the + 3 state becomes more
stable than + 5 state.
Nitrogen can form +1,+2,+3,+4 oxidation states also when it reacts with oxygen. Phosphorous shows
+1,+2,+3,+4 oxidation states in oxoacids.
NH3 N2H4
N2
N2O NO
N2O3
NO2
N2O5
-3
-2
0
+1
+2
+3
+4
+5
(B) By electron sharing.
Elements of Group 15 contain 5 valence electrons (ns2 np3). Out of these five, three electrons are
unpaired electrons (px1, py1, pz1). Thus, these elements can form three covalent bonds by mutual sharing. As
in the case of NH3, PH3, AsH3, SbH3 and BiH3: the two s electrons remain on the atoms of these elements as
lone pair.
If one of the two s-electrons is promoted to a vacant d-orbital, then the total number of unpaired
electrons becomes 5. As a result, an atom of that element can make five covalent bonds.
Since, all other elements (except N) have d-orbitals in their valence shell; hence these elements can
form 5 covalent bonds. For example, phosphorus forms PF5, PCl5.
Phosphorus (P) :
hydrazoic acid
H2P PH2
phosphorus dihydride
The chemical behaviour of nitrogen differs considerably from that of the other elements of this
group. This is due to the following reasons.
Due to p- p overlap, N forms multiple bonds with itself as well as with carbon and oxygen, e.g.,
The tendency to form multiple bonds decreases as we go down the group. It is due to this reason that
elemental nitrogen exists as a diatomic (N 2, N N) molecule. The bond energy
of N N bond is very high (941 kJ mol1).
Due to very high bond energy of N
N bond, nitrogen is almost
non-reactive in its elemental form.
Phosphorus, arsenic, and antimony do not form p- p multiple bonds
hence these elements do not exist as diatomic molecules. To satisfy their
valence shell these elements form tetratomic, tetrahedral molecules such as P 4,
As4 and Sb4. Since,P P single bond is much weaker than triple N N bond,
hence phosphorus is more reactive than nitrogen.
Recently, a few compounds having P = C, P C, P = N, P = P and As = As groups which involve
p- p multiple bonding by phosphorus and arsenic have been synthesized.
Phosphorus and the heavier members of the family do not readily form p- p multiple bonds
whereas multiple bonding of the d- p type can occur readily for these elements.
The d- p bonding is particularly prominent for phosphorus as reflected in the formation of
compounds, such as, 0=PX3 and RN = PX3. In these compounds, the p-orbital of oxygen or nitrogen atom
overlaps sideways with the d-orbital of phosphorus atom to form d- p multiple bond.
Ques.:- Nitrogen can not show higher covalency while other elements of 15th group can, why?
Ans.:- As nitrogen has no d-orbitals in its valence-shell, therefore it can form at the maximum four covalent
bonds, e.g., in NH4+, NR4+ etc. The heavier members of the group can expand their valence shell and form
penta- and hexa-coordinated derivatives such as PCl5, AsF5 and PF6. The higher coordination numbers are
readily adopted when electronegative substituents are present. Since, N cannot expand its valence shell
beyond four hence it does not form pentavalent compounds such as NCl5, NF5 etc.
Ques.:- Nitrogen is very less reactive then other elements of its group, why?
The chemical behaviour of these elements can be described by considering the formation of
various types of compounds as discussed below.
Hydrides of Group 15 Elements:All Group 15 elements give hydrides of the types MH 3 (where M = N, P, As, Sb, Bi), and M 2H4
(where M = N and P).
Phosphine (PH3) and hydrides of other heavier elements of this group are highly poisonous. Some general
characteristics of hydrides of Group 15 elements are given below.
1. Structure. All MH3 hydrides of Group 15 elements are covalent in nature. The central atom (i.e.,
N, P, etc.) are sp3 hybridised. Three hybrid orbitals form covalent bonds with 1s
orbitals of hydrogen atoms, while the fourth orbital is occupied by a lone-pair of
electrons. This gives pyramidal structure to these hydrides. Pyramidal structure
of ammonia (NH3) molecule is shown in Fig. 8.11.
Other MH3 hydrides have similar structure. But, HMH bond angles are
different. For example,
Hydride
:
NH3 PH3 AsH3 SbH3 BiH3
Bond angle :
107 94
92
91
90
The bond angle decreases as the size of the central atom in hydrides increases.
EXPLANATION: In hydrides (MH3 type), the central atom is surrounded by three bond pairs and
one lone pair of electrons. Due to the presence of lone pair of electrons on the central atom, the HMH angle
in MH3 hydrides are less than the tetrahedral angle of 109 28. This can be explained as follows.
The lone pair bond pair repulsions are greater than the bond pair bond pair repulsions. Since, a
bond pair experiences a lesser repulsion from another bond pair than from a lone pair, hence the bond pairs
are brought closer to each other. As a result, the HMH bond angle becomes smaller than the tetrahedral angle
of 109 28. In fact, the HNH bond angle in NH3 is 107.
The strength of lone pair bond pair repulsion depends upon the size of the central atom. As the
atomic size increases, the distance between the bond pairs increases, and therefore, the bond pair bond pair
repulsion decreases. As a result, the lone pair bond pair repulsion will increase with an increase in the
size of the central atom. The increased lone pair bond pair repulsion will push the bond pairs to get closer
to each other, and thereby decrease the bond angle. It is due to this that the HMH bond angle decreases as
we go from NH3 to BiH3.
2. Thermal stability. The stability of these hydrides decreases sharply in going from NH 3 to BiH3.
SbH3 and BiH3 are thermally unstable.
EXPLANATION: As we go down the group, the size of central atom increases therefore the
tendency to form stable covalent bond with small hydrogen atom decreases. As a result the M H bond
strength decreases down the group. As a result, the thermal stability of MH 3 type of hydrides decreases in
going from top to bottom in the group.
3. Basic strength. The basic strength of MH3 type hydrides of Group 15 elements decreases in going
from NH3 to BiH3. Thus, basic strength of MH3 type of hydrides of this Group follows this order,
Except NH3, the hydrides of all other elements are strong reducing agents and react with metal ions,
such as, Ag+ and Cu2+ to produce phosphides, arsenides, and antimonides.
EXPLANATION: Formation of phosphides, arsenides and antimonides involve breaking of MH
bonds. Since, the thermal stability of hydrides of Group 15 elements decreases from NH 3 to BiH3, hence the
reducing character increases in going from NH3 to BiH3.
5. Boiling points. Boiling point of hydrides of Group 15 elements are given below.
Hydride
:
NH3
PH3
AsH3
SbH3
BiH3
Boiling point, K
:
238.5
185.5
210.6
254.6
290
Ques.:- The boiling point of ammonia is unusually higher than that of phosphine. For other
hydrides, the boiling point increases down the group, why?
Ans.:- Unusually high boiling point of ammonia as compared to that of phosphine is due to the
intermolecular hydrogen bonding in ammonia. Phosphine molecules do not form hydrogen bonds with each
other.
The increase in the boiling point of hydrides in
going from PH3 to BiH3 is due to the increased strength of
the Vander Waals forces down the group.
Halides of Group 15 Elements:Elements of Group 15 form two types of halides, viz., trihalides (MX3) and pentahalides (MX5).
All trihalides of N, P, As, Sb and Bi except NI3 are known. Nitrogen does not form pentahalides,
e.g., NF5, NCl5, NBr5 and NI5 due to the absence of d-orbitals in the valence shell of nitrogen. Bismuth
(Bi) also does not form pentahalides. This is because due to inert-pair effect +5 oxidation state of Bi is not
stable.
Some general characteristics of the halides of Group 15 elements are described below.
Trihalides. All the elements of Group 15 form trihalides (except NI3).
(a) All trihalides of Group 15 elements have pyramidal structure in the
gaseous state. In all trihalides, the central atom shows sp3
hybridisation.
(b) Except BiF3, all other trihalides are covalent. The covalent character
of trihalides decreases in going down the group.
EXPLANATION: As we go down in the group, the tendency to form
ionic bonds increases. So, the covalent character decreases.
(c) Trihalides of Group 15 elements (except NF3 and BiF3) get easily
hydrolysed by water.
NCl3 + 3H2O
NH3
+ 3HClO
Oxides of Group 15 Elements:Generally group 15 elements form two types of oxides E 2O3 and E2O5.Nitrogen combine with oxygen
to give a variety of oxides. Typical oxides of group 15 elements are given below.
DINITROGEN (N2)
Preparation of N2 :1. Laboratory method:- In laboratory N2 is prepared by heating aqueous solution ammonium chloride
and sodium nitrite.
NH4Cl (aq) + NaNO2 (aq)
N2 (g) + 2H2O +NaCl (aq)
2. By thermal decomposition of ammonium dichromate:(NH4)2Cr2O7 heat Cr2O3 + N2 (g) + 4H2O
Oxoacids of Group 15 Elements
The acids which contain oxyanions are called oxoacids. Except bismuth, all other elements of Group
15 form oxo-acids.
(a) Oxoacids of nitrogen. Some common, oxoacids of nitrogen are described below. HNO 3 is
stronger than HNO2.
(a) Oxoacids of phosphorus. Phosphorus forms a large number of oxoacids. All these acids are based
on tetrahedral four coordinated phosphorus atom containing at least one P = O unit and one P OH
group. Condensed systems are formed by P O P linkage or P P linkage. Some of the common
oxo-acids of phosphorus are given below.
*pKi = log Ki, where Ki is the acid dissociation constant for the release H+.
Strength of Oxoacids of Group 15 Elements
Following generalizations can help in predicting the strengths of the oxoacids.
(i) The strength and solubility of oxoacids having central atom in the same oxidation state follow the
order, N > P > As > Sb
For example, for an oxidation state of + 5, the strengths of various acids of Group 15 elements
follow the order, HNO3 > H3PO4 > H3AsO4 > H3SbO4
(ii) For the oxoacids involving the same element in different oxidation states, the strength of an acid
depends upon the number of unhydrogenated oxygen atoms attached to the central atom. For example, for
the oxoacids of the type (OH)m ZOn, the acid strength varies directly with the value of n. Thus, nitric acid
(HNO3) is stronger than nitrous acid (HNO2), i.e., HNO3 > HNO2.
The acids H3PO2, H3PO3 and H3PO4 are approximately of equal strength, because all these acids
contain only one unhydrogenated oxygen atom each.
The strength of an acid is described in terms of acid dissociation constant (Ka). The release of each
replaceable hydrogen in the acid molecule is characterised by the corresponding Ka value.
Phosphorus
Atomic no. : 15
Atomic symbol: P
Atomic mass : 31 u
Electronic configuration: 1s2 2s2 2p6 3s2 3p3
Phosphorus is the second member of group 15 of the periodic table and has the electronic
configuration, 1s2 2s2 2p6 3s2 3p3. Although nitrogen and phosphorus belong to the same group, but they
differ from each other considerably in their chemical behaviour.
White phosphorus ignites spontaneously in air. So, phosphorus is always kept under water.
Allotropes of Phosphorus
Phosphorus exists in the following five different allotropic forms.
White or Yellow Phosphorus
White phosphorus is the most common allotropic form of phosphorus. It is obtained by rapidly
cooling the vapours of phosphorus. It is usually prepared either from phosphorite mineral or boneash. When
exposed to light, it becomes yellow.
(i)
White (yellow) phosphorus is extremely reactive.
(ii)
Below 800C, its vapor density corresponds to the formula P4. Above 1700C, it exists as P2.
(iii)
Due to its low ignition temperature (~ 30C), it catches fire spontaneously. Due to this reason, it
is stored under water.
(iv)
It is extremely poisonous.
(v)
It dissolves in carbon disulphide, and turpentine oil.
(vi)
It melts at 44.1C and boils at 280C.
(vii) It glows in the dark due to slow oxidation. This phenomenon is called phosphorescence.
Red Phosphorus
Red phosphorus is stable allotrope at room temperature.
Red phosphorus is formed by heating white phosphorus in the
absence of air at about 250C. It is not poisonous. It is safe to
handle because it does not burn spontaneously at room
temperature.
(i)
Ignition temperature of red phosphorus is high (~
265C).
(ii)
It sublimes on heating. It melts at 610C under
pressure giving yellow liquid, which on cooling gives white phosphorus.
(iii)
It is insoluble in carbon disulphide.
(iv)
Red phosphorus has a polymeric structure, (Fig. 8.19).
Black Phosphorus
Black phosphorus is obtained by heating white phosphorus at about 200C under very higher
pressure (about 4000 atmospheres).
It may also be obtained by heating white phosphorus at 220 370C for 8 days in the presence of
mercury.
(i) Black phosphorus is crystalline in nature, and consists of a layered structure .
(ii) It melts at 587C.It does not burn in air even upto 400C.
(iii) It is a good conductor of electricity.
Scarlet Phosphorus
Comparison of the properties of white and red phosphorus is made in Table 8.12.
Phosphine, (PH3)
Phosphorus forms a number of hydrides. Phosphine (PH3) is the most important hydride of
phosphorus. It was discovered by Gengembre in 1783.
Preparation
(i) In the laboratory, phosphine is prepared by boiling white phosphorus with a concentrated solution
of sodium hydroxide in an inert atmosphere of carbon dioxide, oil gas or hydrogen.
Physical Properties
Some characteristic properties of phosphine are described below.
(i)
It is a colourless gas with disagreeable odour of rotten fish.
(ii)
It is highly poisonous in nature.
(iii)
It is only slightly soluble in water, the solution being neutral towards litmus.
(iv)
It condenses to a colourless liquid (b.p. 188 K), which freezes to a white solid (m.p. 139.5 K).
Chemical Properties
Some important chemical properties of phosphine are given below.
(i) Basic nature. Phosphine is neutral towards litmus. But, it is feebly basic: it is a much weaker base than
ammonia. As a result, it reacts with halogen acids to give phosphonium salts.
For example, with HI, it gives
(ii) With halogens. It burns in chlorine giving phosphorus pentachloride.
(iii)Reducing property. Phosphine acts as a strong reducing agent. It explodes in contact with a very
small amounts of oxidising agent such as nitric acid, chlorine gas etc.
(iv) Formation of coordination compounds. Phosphine does not form any coordination compounds
with the transition-metal cations. However, it forms coordinate compounds with electron deficient
compounds, e.g., with BCl3 it gives,
Uses of Phosphine
(i)
It is used for producing smoke screens. Calcium phosphide is used in smoke screens. Phosphine
obtained catches fire to give the needed smoke.
(ii)
In Holmes signals. A mixture of calcium carbide and calcium phosphide is taken in a container
which is pierced and thrown into the sea. Phosphine liberated catches fire and lights up acetylene.
Burning gases serve as a signal to the approaching ships.
Group 16 Elements-Introduction
The elements Oxygen (O), Sulphur (S), Selenium (Se), Tellurium (Te) and Polonium (Po) constitute
group 16 of the periodic table. The first four members of this group are non-metallic in nature. These are
called chalcogens (means ore-forming) because most of these metals occur in nature as oxides, or
sulphides. The last member of this family, i.e., polonium, is a radioactive element with a very short half-life.
The elements of group 15 have their p-orbitals in the valence shell, half-filled. Half-filled
configurations are more stable. Due to the extra-stability of the half-filled configurations more energy is
required to remove an electron from the outermost shell for the elements of group 15.
(vi) Oxidation states
The outer electronic configuration of group 16 elements can be described as ns2np4. Being strongly
electronegative, these elements complete their outer shells by gaining two electrons. Thus, all the elements
of group 16 show an oxidation state of 2. However, these elements also show other oxidation states as
follows. Oxygen shows an oxidation state of + 2 in F 2O, and 1 in peroxides (O2 2). Other elements of
group 16 exhibit oxidation states of + 2, + 4 and + 6 also. The oxidation states of + 4 and + 6 being more
stable.
For sulphur, selenium and tellurium, the oxidation states of + 4 and + 6 are important. The + 4 state
is more stable for Se, Te and Po, than + 6 state. This is due to the availability of d-orbitals in the valence
shells of the atoms of these elements. For sulphur the scheme is shown below.
Sulphur forms numerous allotropes which may be grouped into three classes:
(a) homocyclic species containing 6 to 20 sulphur atoms of which the commonest and
thermodynamically the most stable form is the orthorhombic S8 molecule.
(b) various chain polymers known as catena-Sn sulphur,
(c) unstable small molecules Sn(n = 2 to 5) which exist in various concentrations in liquid sulphur at
higher temperatures and in sulphur vapour. At approximately 1000 K, S2 species predominates.
Like oxygen (O2), S2 is paramagnetic.
Selenium exists in six allotropic modifications: of these three are red monoclinic forms containing
Se8 rings. The stable form is grey hexagonal monoclinic allotrope which consists of infinite spiral chains and
does not have a sulphur analogue.
To break hydrogen bonds, energy is needed. At the boiling point, all the remaining hydrogen bonds
in water must be broken. This would need some more energy. As a result, the boiling point of water is
higher. Thus, the existence of hydrogen bonds between water molecules is mainly responsible for its
abnormally high melting and boiling points, (or for its low volatility).
Hydrides of other elements of this group have the van der Waals forces. Thus, it is because of the
absence of hydrogen bonding that the hydrides of sulphur, selenium, and tellurium are highly volatile and
gaseous.
The strength of the van der Waals' forces increases with molecular mass. So, as we go down the
group, the strength of the van der Waals forces increases. Therefore, the boiling point of hydrides increases
and the volatility decreases when we go from H2S to H2Te.
(iii) Thermal stability. The thermal stability of hydrides of Group 16 elements in the order,
H2O > H2S > H2Se > H2Te > H2Po
EXPLANATION: The M H bond energy and hence thermal stability decreases as we go down the
group from O to Po due to an increase in the size of the central atom (Table 8.14). Water molecule is
however, more stabilized due to hydrogen bonding.
(v) Acidic nature. The hydrides of group 16 elements are weakly acidic. The acidic character of
these hydrides increases with increasing atomic number. Thus, the acid strength of these hydrides increases
as follows.
All these hydrides, viz., H2S, H2Se and H2Te are diprotic acids.
EXPLANATION: In all these hydrides, the M H bonds are polar. As a result, these hydrides ionise
in aqueous solutions to give free H3O+. Thus, these hydrides act as acids.
The atomic size increases in going from O to Te. As a result, the energy of the M H bond decreases,
and the bond breaking process becomes easier. This leads to an increase in the acid strength of hydrides in
going down the group.
(vi) Reducing character. All hydrides of Group 16 elements except H2O, are reducing agents. The
reducing power of these hydrides increases in going from H2S to H2Te.
Halides of Group 16 Elements
Sulphur forms halides of the type, S 2X2 (X = F, Cl, Br, I), SX 2 (X = F, Cl, Br), SX 4 (X = F, Cl) and SF6.
Selenium and tellurium form hexafluorides. Fluorides, chlorides and bromides of selenium and tellurium in
the oxidation states of + 1, + 2 and + 4 are also known.Certain important features of halides of group 16
elements are described below.
(i) The stability and variety of halogen compounds formed by the elements of Group 16 decrease
with the increasing atomic number of the halogen, i.e., in the order
Fluoride
>
Chloride
>
Bromide
>
Iodide
This means, fluorides are the most stable, while iodides are the least stable of the halides of
group 16 elements.
(ii) Sulphur, selenium and tellurium show maximum valency of six only in their fluorides. It is due to
the small size and the most electronegative nature of fluorine atom. The central atom in hexafluorides shows
sp3d2 hybridisation. So, hexafluorides have octahedral structures.
(iii) The compounds of fluorine with oxygen are called oxygen fluorides. For example, the compound
F2O is actually written as OF2 and is named as oxygen difluoride. This is because fluorine is more
electronegative than oxygen.
The compounds of chlorine, bromine and iodine
with oxygen
are called halogen oxides, because oxygen is more
electronegative than chlorine, bromine and iodine. For
example, ClO2
is named as chlorine dioxide, Cl2O7 is named as chlorine
heptoxide etc.
(ii)
Oxides of sulphur are more stable than those of selenium, i.e., SO2 is more stable than SeO2, SO3 is
more stable than SeO3.
(iii)
Dioxides. Sulphur (S), selenium (Se) and tellurium (Te) burn in air to form
dioxides of the type MO2. All dioxides are acidic in nature. The acidic
character decreases as we go down the group.
Structures of dioxides.
SO2 being a discrete molecule exist in the gaseous state. Hybridisation of S
in SO2 is sp2. The OSO bond angle is 119.5. Structure of SO 2 in the gaseous state
is shown in Fig. 8.25.
Selenium and tellurium dioxides are solids having polymeric chain or layer
structure. Structure of SeO2 in the solid state is given in Fig. 8.26. It consists of a
zig-zag chain. Hybridisation of Se in SeO2 is sp3.
TeO2 and PoO2 are crystalline ionic solids.
EXPLANATION: As compared to sulphur, the selenium and tellurium oxides have higher lattice
energy. Therefore, in the case of Se and Te, chain-structure is more stable than p d bonding. Therefore,
SeO2 and TeO2 are solids.
(iv) Trioxides. All the group 16 elements form trioxides, MO3. Sulphur trioxide in the gas phase
exists as planar triangular molecular species (Fig. 8.27), although in the solid state it can exist as a linear
chain or a cyclic trimer.
Selenium trioxide (SeO3) solid is a cyclic tetramer (Se4O12) as shown in Fig. 8.27. TeO3 is a solid
with a network structure in which TeO 6 octahedra share all vertices. SO2 and SO3 are the most important
oxides from industrial point of view.
Oxoacids of Group 16 Elements
Sulphur forms many oxoacids:
sulphuric acid being the most important.
The principal acids formed by sulphur
are shown in Fig. 8.28. Some of the
acids of sulphur, e.g., sulphurous and
dithionic acids are known in the form of
their salts only, as these acids in the free
state are unstable and cannot be
isolated.
Structure of Sulphur
Both rhombic and monoclinic sulphur
consist of eight
sulphur atoms combined in a ring but differ from
each other in the
symmetry of their crystals. The S8 ring is
puckered
as
shown in Fig. 8.32. Rhombic sulphur is stable at
room
temperature whilst monoclinic sulphur is stable
above 369 K,
(96C).
At higher temperatures, open chains
consisting
of
sulphur atoms covalently bonded to each other are formed. Above 1000 K, it exists as S2 molecule.
(ii) By roasting sulphides. Zinc sulphide, iron pyrites, etc., when roasted in the air produce SO 2.
This method also is used for preparing large volumes of sulphur dioxide needed in industry, e.g., in the
manufacture of sulphuric acid.
When the colourless product is allowed to stand in the air, it is reoxidised by the atmospheric oxygen
to its original colour. So, bleaching with sulphur dioxide is temporary.
Colourless vegetable product
+
[O]
The bleaching action of chlorine, however, is due to its oxidising nature, and is permanent.
Uses of SO2 Gas
(i)
For the manufacture of sulphuric acid.For bleaching hair, silk, wool, fur etc.
(ii)
For refining and decolourising cane juice in sugar factories.
(iii)
For the preparation of sulphites.
(ii) By dehydration of sulphuric acid. When concentrated sulphuric acid is heated with P4O10,
sulphur trioxide is formed.
Physical Properties
Sulphur trioxide exists in three forms.
(a) -SO3. It forms long, transparent, ice-like needles. It melts at 290 K and boils at 318 K.
(b) -SO3. The -SO3 gets gradually converted to -SO 3 below 298 K. It form a silky, asbestos like
needles and melts at 305.5 K.
(c) -SO3. When -SO3 is completely dried, -SO3 is obtained. It resembles -SO3 in appearance but
sublimes at ordinary temperatures. It melts at 335 K under a pressure of 2.5 atm.
Chemical Properties
Theory
The chemistry involved in the contact process is described below.
(i) Production of SO2. Sulphur dioxide (SO2) is obtained by burning sulphur or iron pyrites.
(ii) Catalytic oxidation of SO2 to SO3. Sulphur dioxide is oxidised by air in the presence of a
catalyst V2O5 to give sulphur trioxide.
The reaction is exothermic, and proceed with a decrease in volume. Therefore, according to
LeChateliers principle, the favourable conditions for the maximum yield of sulphur trioxide are:
(a) Air or oxygen required for the oxidation of sulphur dioxide must be in excess.
(b) The temperature must be low. A temperature between 350 450C gives the maximum yield of the
product.
(c) The pressure must be high, about 2 atmospheres.
(d) Platinised asbestos was used as a catalyst previously, but now-a-days it is replaced by much cheaper
vanadium pentoxide (V2O5). Vanadium pentoxide remains uneffected by the impurities, while
platinised asbestos is poisoned by the impurities in the gases.
(e) The gases used (SO2 and O2) must be free of impurities, viz., dust particles, arsenious oxide etc., to
prevent catalyst poisoning.
(iii) Conversion of sulphur trioxide into sulphuric acid. Sulphur trioxide is dissolved in concentrated
sulphuric acid to produce oleum or fuming sulphuric acid.
(iv) Conversion of oleum to sulphuric acid. Oleum is diluted with a calculated amount of water to get
sulphuric acid of desired concentration.
(i) Acidic character. Sulphuric acid is a strong dibasic acid and ionises in the solution in two steps:
Therefore, sulphuric acid forms two sets of salts normal sulphates (such as sodium sulphate and
copper sulphate) and acid sulphates (e.g., sodium bisulphate).
Group 17 Element-Introduction
Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I) and Astatine (At) are the members of the
halogen family. This family constitutes group 17 of the periodic table. The word halogens is derived from
the Greek words halo (sea salt) and gen (producer). Astatine, the last member of the series is radioactive.
The isotopes 211At (half-life 7.2 hours) and 210At (half-life 8.3 hours) are short-lived. Therefore, not much is
known about the chemistry of astatine.
the most electronegative shows only 1 oxidation state. Other halogens show both the negative as well as
positive oxidation states. The common oxidation states shown by the halogens are,
Fluorine
1
Chlorine
1, + 1, + 3, + 5, + 7
Bromine
1, + 1, + 3, + 5
Iodine
1, + 1, + 3, + 5, + 7
Bromine and chlorine also exhibit the oxidation states of + 4 and + 6 in oxoacids and oxides.
EXPLANATION: Fluorine does not show higher/variable oxidation states because it does not have
d-orbitals in its valence shell.
Chlorine, bromine and iodine have d-orbitals in their valence shell. Due to the availability of the dorbitals in higher halogens, the paired electrons in the outer s- and p-orbitals can be unpaired by promoting
them to the vacant d-orbitals. This open up the possibilities of variable oxidation states as described below.
EXPLANATION: The hydration enthalpy depends inversely on the ionic radius. In going from down
the group, the ionic radii increases. So, the hydration enthalpy becomes smaller.
(vi) Colour All halogens are coloured, viz.,
Fluorine
Chlorine
Bromine
Iodine
Light yellow Greenish yellow
Reddish brown
Violet (purple)
Thus, the colour of the halogens gets darkened as one goes from F to I.
EXPLANATION. When molecules or atoms absorb energy (may be in the form of light), its outer
electrons get excited to higher energy levels. When atoms / molecules absorb light in the visible region, the
complementary colour is emitted out. Halogens absorb light in the visible region. Fluorine absorbs blue light
and emit pale-yellow light. So, fluorine appears light yellow in colour. Iodine absorbs the yellow light, so it
appears violet in the visible light.
2F + X2
(X = Cl, Br, I)
Here, fluorine oxidizes Cl, Br and I to chlorine, bromine and iodine.Chlorine oxidizes Br and I
to bromine and iodine.Bromine oxidizes I to iodine.
Due to their high reduction potentials (E), fluorine, chlorine and bromine decompose water to
liberate oxygen.
X2 X = F, Cl, Br) + H2O
HOX + HX
HOX
O + HX
O+O
O2(g)
Halides of Hydrogen: Hydrides of Group 17 Elements
Halogens combine with hydrogen to form hydrogen halides. The reactivity of halogens decreases in
going from fluorine to chlorine. Thus,
F2
>
Cl2
>
Br2
>
I2
most reactive least reactive
(i) Physical state. Hydrogen fluoride (HF) is a low boiling liquid (boiling point: 292 K), while HCl,
HBr and HI are gases at room temperature.
Relatively higher boiling point of hydrogen fluoride (HF) is due to hydrogen bonding in hydrogen
fluoride. Hydrogen bonding in hydrogen fluoride leads to an associated structure of hydrogen fluoride, i.e., it
may exist as (HF)n.
......... H+ F ......... H+ F ......... H+ F
(ii) Nature. Hydrogen halides are polar covalent diatomic molecules in the gaseous phase. H X
bond in halides has an appreciable ionic character. Hydrogen halides dissolve readily in water forming
hydrohalic acids. These hydrohalic acids ionise in solution to form H3O+ and X ions.
(iii) Acidic nature. In aqueous solutions, hydrogen halides behave as acids. This is due to the
formation of H3O+ ions in solution.
The tendency of a hydrogen halide to give H 3O+ in solution is described in terms of its acid strength.
The acid strength of hydrogen halides in water follows the order.
HI
>
HBr >
HCl >
HF
Thus, HI is the strongest and HF is the weakest acid in this series.
EXPLANATION. The strength of any hydrohalic acid is determined by the tendency of the reaction,
H2O + H X
H3O+ + X
in the forward direction. The tendency of this reaction in the forward direction in governed by the energy
required to break the H X bond. So, lesser the bond energy of the H X bond, stronger is the acid.
The bond energy decreases in going from HF to HI. Therefore, H I bond requires the least and H
F bond requires the maximum energy in this series of halides. As a result, HI is the strongest and HF is the
weakest of the hydrohalic acids. Thus, the acid strength of hydrohalic acids follows the order
HI
>
HBr >
HCl >
HF
(iv) Thermal stability. HF is the most stable, while HI is the least stable hydrogen halide. Thus, the
thermal stability of hydrogen halides follow the order,
HF
>
HCl >
HBr >
HI
thermally most stable thermally least stable
(v) Reducing character. HF is not a reducing agent, while HI is a strong reducing agent. Thus, the
reducing character increases in going from HF to HI. This because
(a) the bond energy of H I bond is low
(b) I has a very strong tendency to get oxidized to I2. As a result, it acts as a strong reducing agent.
Oxides of Group 17 Elements
Fluorine forms two oxides, F2O and F2O2, which are properly called oxygen fluorides because of the
higher electronegativity of fluorine. Chlorine, bromine and iodine form oxides in which the oxidation state
of the halogen ranges from + 1 to + 7. The oxides of halogens are listed in Table 8.21.
Hypohalous Acids
Preparation. Hypohalous acids are formed by the disproportionation of halogens in water,
X2 + H2O
HOX + HX
hypohalous acid
non-linear
acids give
NaClO is
cold NaOH.
sodium hypochlorite
In these molecules, apart from OH group, there are differing numbers of oxygen atoms on the
central atom . Due to the electron-withdrawing tendency of the unhydrogenated oxygen atoms, Cl atom
developes a slight positive charge (it is called formal charge)*. This leads to shifting of electron-density
away from the oxygen atom of the OH group. As a result, release of H atom as a proton (H +) becomes
easier. Thus, the acid having greater number of unhydrogenated O atoms on the central atom tends to
be stronger.
Available Chlorine in Bleaching Powder
A good bleaching powder contains about 35% of chlorine. This is called available chlorine. The
available chlorine is defined as the mass of the chlorine gas liberated when 100 g of bleaching powder is
made to react with an excess of dilute acid.
Interhalogen Compounds
Compounds containing two or more halogens are called interhalogen compounds. Halogen
combine with each other due to electronegativity difference between them. There are three different types of
interhalogen compounds,
(i) Molecular interhalogen compounds.
(ii) Compounds containing polyhalogen anions of the type XY2n , where n = 1, 2, 3, ...... For example,
ICl2, ICl4, I3.
(iii) Compounds containing polyhalogen cations of the type, XY2n+ , where n = 1, 2, 3, ...... Forexample,
ICl2+, ICl4+.
Molecular Interhalogen Compounds
There are four different types of molecular interhalogen compounds.
These are illustrated below.
Group 18 Element-Introduction
Group 18 of the periodic table accomodates six elements. These are, Helium (He), Neon (Ne), Argon
(Ar), Krypton (Kr), Xenon (Xe) and Radon (Rn). The last member, radon is a radioactive element.
These gases were earlier called inert gases, because of their chemical inertness (non-reactivity under
normal conditions). This chemical inertness was attributed to their stable electronic configurations. In the
last four decades, a good number of compounds of these elements have been prepared, thereby discarding
the earlier belief of their inertness. These gases, now-a-days, are called noble gases (the gases with a little
chemical reactivity). Since, these gases are present in or around the earth in very small amounts, these gases
are also called rare gases.
Uses of Neon
It has moderately high electric discharge potential. It has beautiful orange-red discharge colour and is
filled into electric discharge tubes used in decorations and for advertisements. The colour so obtained in a
discharge tube depends upon.
(i)
nature of gas or vapour mixed with neon.
(ii)
the type of glass used in the preparation of discharge tube.
Some typical colours given by neon-mixtures are given below.
Uses of Argon
(i)
It is used in gas-filled electric lamps.
(ii)
A mixture of argon and mercury vapour is used in fluorescent tubes..
Uses of Krypton
(i)
It is used for filling luminous sign tubes.
(ii)
It is used in filament lamps.
Uses of Xenon
(i)
It is used for filling radio and television tubes.
Uses of Radon
(i)
In the preparation of ointment for the treatment of cancer and other diseases.
(ii)
Used in the scientific research on radioactivity
Compounds of Xenon
Xenon in its compounds exhibits positive oxidation states from +2 to +8. For example,
Compound
:
XeF2 XeF4 XeF6 XeO4
Oxidation state of Xe :
+2
+4 +6
+8
Here, we would discuss some fluoro and oxyfluoro compounds of xenon.
Xenon difluoride, (XeF2)
In this compound, Xe appears in + 2 oxidation state.
Preparation. It is prepared in the laboratory by heating xenon and fluorine in the molar ratio 1 : 3 in
a nickel vessel at 400C. The reaction products are quenched at 50C. XeF 2 is isolated by vacuum
sublimation.
It can also be obtained by reacting together Xe and O2F2 (oxygen difluoride) at 118C.
Structure and bonding. XeF2 is a linear molecule, F Xe F. Valence bond repre-sentation of XeF 2 may
be explained, if one of the 5p electrons is promoted to the 5d orbital.
(i)
One 5s, three 5p and two 5d atomic orbitals of xenon hybridize to give six sp3d2 hybridized orbitals.
The four singly occupied hybridized orbitals are used by four fluorine atoms for bond formation and the rest
two are occupied by lone pairs as shown in Fig. 8.45.
XeO3 + 6HF
Structure of XeF6
The Lewis structure of XeF6 is given in Fig. 8.46.
Xenon trioxide, (XeO3)
(i) By the hydrolysis of xenon fluorides. XeF4 and XeF6 on hydrolysis gives XeO3.
XeF6 + 3H2O
XeO3 + 6HF
(i)
6XeF4 + 12H2O
XeO3 + H2O
.
H2XeO4 l H+ + HXeO4
xenic acid
Structure of XeO3
In XeO3, Xe exhibits sp3 hybridisation. One of the hybrid orbitals is occupied by a lone pair, while
other three are involved in the bond formation. This leads to a trigonal pyramidal structure of XeO3.
Xenon oxytetrafluoride, (XeOF4)
Preparation.
Xenon oxytetrafluoride (XeOF4) can be obtained by partial hydrolysis of XeF6.
XeF6 + H2O
XeOF4
+
2HF
Xenon
oxytetrafluoride
Structure.
In this compound, xenon exhibits + 6 oxidation state and sp3d2 hybridisation. This gives square
pyramidal structure for XeOF4.