Chemistry of the Elements
4
GROUP SEVEN ELEMENTS
4.1
PHYSICAL PROPERTIES & TRENDS
Formula
Electronic Structure
Electronegativity
Bond Energy/ KJ mol-1
M.P/0C
B.P/0C
Ionization Energy/KJ mol-1
Electron Affinities/KJ mol-1
Atomic Radius/PM
NM
Ionic Radius/PM
Eox2/x-/V
Oxidising Power
Ox. No.
Standard Enthalpy / ΔHDo / KJ mol-1
Colour/State
At No.
Electronic Config.
Fluorine
[He] 2s22p5
4.0
158
-220
-188
1681
361
64
0.072
133
+2.87
MOST
-1
79.1
G (Pale yellow)
9
2.7
Chlorine
[Ne] 3s23p5
3.0
242
-101
-34
1231
388
99
0.099
181
+1.36
Bromine
[Ar] 3d104s24p5
2.8
193
-7
58
1140
365
111
0.114
196
+1.09
-1, +1, +5, +7
122
G, (Pale green)
17
287
-1, +1, +5, +7
111
L, (Red brown)
35
28187
Iodine
[Kr]4d105s2sp5
2.5
151
114
183
1010
332
128
0.133
219
+0.54
LEAST
-1, +1, +5, +7
106
S, (Black)
53
281818
TABLE 4.1 – Physical Properities Of The Transition Elements (First Row)
All the halogens exist as diatomic molecules. F2, Cl2, Br2, I2. The two atoms are linked by
covalent bonds. Fluorine and chlorine are gases, bromine is a liquid and iodine is a solid. Fluorine
is pale yellow, chlorine is pale yellow, bromine is red brown and iodine is shiny black. The
decreasing volatility is related to the increasing Van der Waals forces with increasing relative
molecular mass. This results in increasing melting points and boiling points and molar enthalpy
changes of vaporization. The halogens are oxidizing agents. They are non – polar simple
molecules and are more soluble in organic solvents, the colour of the solution depending upon the
particular halogen and the solvent. In tetrachloromethane and cyclohexane, chlorine is colourless,
bromine is read and iodine is violet. In polar (electron donating) solvents such as water, ethanol
and propanone (acetone) bromine and particularly iodine tends to give brownish solutions. The
change in colour is due to the formation of complexes invoking charge transfer. This involves a
donation of electrons from the polar solvent molecules to the iodine molecule. It is thought that
the electron normally located in an orbital of the solvent molecule can absorb energy from visible
light and jump into an orbital of the iodine molecules. Thus the complex resumes visible light of a
particular wavelength and appears coloured.
The blue compound which iodine forms with starch is also a complex. The iodine molecule is
absorbed reversibly within the helical structure chain of glucose molecule which makes up starch.
Copyright © Pooran Appadu
Chemistry of the Elements
4.2
HALOGENS AS OXIDIZING AGENTS
The halogen accepts electrons during reactions and acts as oxidizing agents.
Hal2
+ 2e-
2Hal-
Fluorine is the most reactive halogen and the most powerful oxidizing agents. The order of
decreasing power of oxidizing agents is:
F2 > Cl2 > Br2 > I2
The electrode potentials (hal2/Hal-) for the halogens are:
Eo
= F2 / F- = 2.87 V
F2 (g) + 2e-  2F- (aq)
;
Eo
= +2.87
Cl2 (g) + 2e- 
2Cl- (aq)
;
Eo
= +1.36
Br2 (g) + 2e- 
2Br- (aq)
;
Eo
= +1.09
;
Eo
= +1.09
2Fe2+ (aq)
;
Eo
= +0.77
S2O62- + 2e-  2S2O32-
;
Eo
= +1.09
I2 (g) + 2e- 
Fe3+ (g) + e- 
2I- (aq)
F2, Br2 and Cl2 will all oxize Fe2+ (aq) to Fe3+ s
I2, however, cannot remove electrons from iron (III) ions to form iron (II) ions.
Fe2+
Fe2+
Eo =
-0.77
Br2
2Br-
Eo =
1.09
= +0.32
The reaction is feasibile as Eo = positive.
Fe2+
Fe2+
Eo =
-0.77
I2
2I-
Eo =
0.54
= -0.23
The reaction is not feasible as Eo = negative.
Reactivity is largely a matter of rates of reaction, which in turn is related to the activation energy
of the reactions. Fluorine has the lowest bond energy. A collision involving a fluorine molecule
needs less energy to break the molecule apart than does the same collision with, say, a chlorine
molecule. As a result fluorine reacts more readily than chlorine.
Copyright © Pooran Appadu
Chemistry of the Elements
Reactivity is in the order:
F2 > Cl2 > Br2 > I2
Fluorine has only one oxidation i.e. -1. It is restricted to the n = 2 shell.
2p
2s
1s
The other elements have empty d orbitals which allow promotion of the electrons from p orbitals
to d orbitals e.g.
3d
3p
3s
Ground State
4.3
Executed State
REACTIONS WITH HYDROGEN
F2
Reacts explosively even in dark at -200 oC.
Cl2
Reacts explosively in sunlight, slowly in dark below 200 oC.
Br2
Reacts above 200 oC and at lower temperature with Pt catalyst.
I2
Reacts to form equilibrium mixtures of H2, I2 and HI.
OTHER HALOGENS REACTIONS
The oxidizing power of the halogen is measured by the Eo values stated earlier e.g.
Cl2
+ 2e-  2Cl-
Eo
= +1.36
Br2
+ 2e-  2Br-
Eo
= +1.07
Cl2
+ 2Br-
Eo
= +0.29
Br2 + 2Cl-
Chlorine will oxidize bromide and iodide ions bromine will oxidize iodide ions. Fluroine will
oxidize chloride, bromide and iodide ions.
IODINE AND SODIUM THIOSULPHATE
Copyright © Pooran Appadu
Chemistry of the Elements
2S2O32- (aq)
S2O2-6 (aq) + 2e
Eo
=
I2(aq) + 2e-
2I- (aq)
Eo
= +0.54
Eo
= +0.4605
2S2O32- (aq)
S2O62- (aq) + 2I- (aq)
+ I2 (aq)
-0.0895
REACTION WITH WATER
The standard electrode potential for oxygen is Eo = +1.23 V
O2 (g) + 4H+ (aq) + 4e-  2H2O (l)
Eo
= +1.23 V
F2 and Cl2 are capable of oxidizing water while bromine and iodine are not. Chlorine reacts with
water slowly to form HCl and HClO (chloric (I) acid)
Cl2 (g) + H2O (l) 
HCl (aq) + H+ (aq) + Cl- (aq)
Chloric (I) acid decomposes to give oxygen slowly (sunlight accelerates the decomposition)
2H+ (aq) + 2Cl- (aq) + O2 (g)
2HClO (aq)
In the presence of a reducing agent chloric (I) and acts as an oxidizing agent.
HClO (aq) + H+ (aq) +
2e-
Cl- (aq) + H2O (l)
REACTIONS OF SOLID HALIDE WITH CONCENTRATED SULPHURIC ACID
(H2SO4)
Chloride
Bromide
Cl2 (g)
+
H2SO4
+ MnO2
Br2 (g)
Iodide
4.4
I2 (g)
HYDROGEN HALIDES
ΔHo+ /KJmol-1
ΔHoB.D / KJ mol-1
B.P/0C
pKa
HF
-270
+560
20
3.25
HCl
-92
+430
-85
-7.4
HBr
-36
+470
-67
-9.5
HI
+26
+300
-35
-10
TABLE 4.2 – Physical Properities Of The Hydrogen Halides
Thermal stability decrease in the same order with their compounds and acid strength also
decrease in that order.
HF < HCl < HBr < HI
Copyright © Pooran Appadu
Chemistry of the Elements
4.5
HALIDE REACTIONS
REACTIONS OF HALIDE ION WITH AGNO3 AND NH3
Chloride
Bromide
Iodide
Addition of AgNO3 in the presence of HNO3
White ppt. of silver chloride>
Ag+(aq) + Cl-(aq)  AgCl(s)
Pale yellow ppt. of silver bromide
Ag+(aq) + Br-(aq)  AgBr(s)
Yellow ppt. of silver iodide.
Ag+(aq) + I-(aq)  AgI(s)
Addition of NH3 solution
Ppt. dissolves clean solu.
Ppt. partly dissolves.
Ppt. does not dissolve.
TABLE 4.3 – Reactions Of Haliod Ions With Ag
REACTION WITH CONCENTRATED H2SO4
Chloride
Bromide
HCl (g)
+
H2SO4
HBr (g) + Br (g)
Iodide
I2 (g)
REACTIONS OF CHLORINE WITH COLD AND HOT NAOH
(aq)
There are two types of change depending on the temperature of the alkali.
(i)
Cold Dilute Alkali
Cl2 (g) + 2NaOH (aq)
(ii)
NaCl2 (aq)
+ NaClO (aq) + H2O (l)
Hot Concentrated Alkali
3Cl2 (g) +
6OH- (aq)
5Cl- (aq) + ClO3- (aq) + 3H2O (l)
The main difference between the two is that, with cold dilute alkali, chlorate (I), ClO -, ions are
the products, while with hot concentrated alkali, chlorate (V), ClO3- ions are made.
N.B. The chlorate (I) that is formed may decompose to form a chloride and a chlorate (V).
3ClO- (aq)
2Cl- (aq) + ClO3- (aq)
The decomposition is slow at room temperature but fast at 70 oC. A reaction like this in which a
species is simultaneously oxidized and reduced is called disproportionation reaction. Part of the
ClO- is oxidized to ClO3- whiles the rest is reduced to Cl-.
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