Transcript Halogens - ilc.edu.hk

Characteristic
Properties of the
Halogens
1
Introduction
• Group VIIA elements include
 fluorine
 chlorine
 bromine
halogens
 iodine
(Salt producers)
 astatine
2
Introduction
•
Astatine
 chemistry not much known
 radioactive
 the total amount present in the
Earth's crust is probably less than
30 g at any one time.
3
Halogens are p-block elements
•
 outermost shell electronic configuration
of ns2np5
4
Halogens are p-block elements
•
 one electron short of the octet structure
5
Introduction
•
In the free elemental state
they complete their octets by sharing
their single unpaired p-electrons
6
When halogens react with other elements
they either
gain an additional electron to form halide ions
or
share their single unpaired p-electrons to form
single covalent bonds
7
High Electronegativity / Electron Affinity
 highest among the elements in the same period
 have a high tendency to attract electrons
 strong oxidizing agents
8
High Electronegativity / Electron Affinity
 -1 is the most common oxidation state of
halogens in their compounds
9
Ionic :
NaF, NaCl, NaBr, NaI
Covalent :
HF, HCl, HBr, HI
Variable Oxidation State
All halogens (except fluorine) can expand their
octet of electrons by utilizing the
vacant,
energetically low-lying
d-orbitals.
10
11
“Electrons-in-boxes” diagrams of the
electronic configuration of a halogen atom of
the ground state and various excited states
+3
+1
O
H
H
Cl
Cl
O
O
The half-filled orbital(s) overlap(s) with those of
more electronegative atoms (e.g. O)
 positive oxidation state (+1, +3, +5, +7)
O
O
H
O
12
Cl
+5
+7 Cl
H
O
O
O
O
Various oxidation states of halogens in their ions or compounds
Oxidation state
of halogen
–1
Ion / Compound
F–
Cl–
Br–
I–
HF
HCl
HBr
HI
Cl2
Br2
I2
OF2
0
+1
+3
13
F2
Cl2O
Br2O
HOCl
HOBr
OCl–
OBr–
HClO2
ClO2–
Various oxidation states of halogens in their ions or compounds
Oxidation state
of halogen
+4
+5
Ion / Compound
ClO2
BrO2
HClO3
HBrO3
I 2O 5
ClO3–
BrO3–
HIO3
IO3–
+6
+7
14
Cl2O6
BrO3
Cl2O7
H5IO6
HClO4
HIO4
ClO4–
IO4–
•
Fluorine (1)
 the most electronegative element
 only one unpaired p electron
available for bonding
 oxidation state is limited to –1
15
• Fluorine (1)
 cannot expand its octet
 no low-lying empty d orbitals
available
 the energy required to promote
electrons into the third quantum
shell is very high
Absence of HFO, HFO2, HFO3, HFO4
16
Variation in Physical Properties
1. Melting point / boiling point  down the
group
17
Halogen
Melting point (C)
Boiling point (C)
Fluorine
–220
–188
Chlorine
–101
–34.7
Bromine
–7.2
58.8
Iodine
114
184
Astatine
302
380
Variations in melting point and
boiling point of the halogens
18
Variation in Physical Properties
1. Melting point / boiling point  down the
group
The molecular size  down the group
 The electron cloud is more easily polarized
 Induced dipoles are formed more easily
 Stronger London dispersion forces
19
2. Colour becomes darker down the group
Halogen
Colour
20
F2
Pale
yellow
Cl2
Br2
Greenish Reddish
yellow
brown
I2
Violet
black
chlorine
Appearances of halogens at room
temperature and pressure: chlorine
21
bromine
Appearances of halogens at room
temperature and pressure: bromine
22
iodine
Appearances of halogens at room
temperature and pressure: iodine
23
Colour
• All halogens
 coloured
 the absorption of radiation in the
visible light region of the
electromagnetic spectrum
The colour is due to the unabsorbed
radiation in the visible light region
24
Colour
• Fluorine atom
 has the smallest size
 absorbs the radiation of relatively
high frequency (i.e. blue light)
 appears yellow (the unabsorbed
radiation)
25
Colour
• Atoms of other halogens
 larger sizes
 absorb radiation of lower frequency
26
Colour
• Iodine
 absorbs the radiation of relatively
low frequency (i.e. yellow light)
 appears violet
27
Q.1 The colour of astatine is black.
28
Colour
• Halogens
 different colours when dissolved in
different solvents
29
Colours of halogens in pure form and in solutions
Colour
Halogen
in pure form
in water
in 1,1,1-trichloroethane
F2
Pale yellow
Pale yellow
Pale yellow
Cl2
Greenish yellow
Pale yellow
Yellow
Br2
Reddish brown
Yellow
Orange
Violet black
Yellow (only
slightly soluble)
Violet
I2
Brown in KI(aq)
30
Colour
• Halogens
 non-polar molecules
 not very soluble in polar solvents
(such as water)
 but very soluble in organic solvents
(such as 1,1,1-trichloroethane)
31
(a)
(b)
(c)
Colours of halogens in water:
(a) chlorine; (b) bromine; (c) iodine
32
(a)
(b)
(c)
Colours of halogens in 1,1,1-trichloroethane:
(a) chlorine; (b) bromine; (c) iodine
33
3. Electron Affinity
 down the group
Halogen
E.A.
kJ/mol1
34
F
Cl
Br
I
At
-322
-349
-335
-295
-270
The number of electron shells and size of atoms 
down the group

The nuclear attraction for the additional
electron  down the group

Electron affinity  from Cl to I
35
Atoms of fluorine have the smallest size among the
halogens

The addition of an extra electron to the small
quantum shell(n=2) results in great repulsion
among the electrons.

Fluorine has a lower electron affinity than Cl
and Br.
36
4. Electronegativity
 down the group
37
Halogen
F
Cl
Br
I
At
Electronegativity
4.0
3.0
2.8
2.5
2.2
The number of electron shells and size of atoms 
down the group

The nuclear attraction for the bonding
electrons  down the group

Electronegativity  down the group
38
Fluorine has the highest electronegativity because it
is the most reactive elements.
The electronegativity of fluorine is arbitrarily
assigned as 4.0.
39
Variation in Chemical Properties
Reactivity : F2 > Cl2 > Br2 > I2
React by gaining electrons
Oxidizing power : F2 > Cl2 > Br2 > I2
40
1. Reactions with Sodium
• All halogens
 combine directly with sodium to
form sodium halides
 the reactivity decreases down the
group from fluorine to iodine
41
1. Reactions with Sodium
• Fluorine
 react explosively to form sodium fluoride
2Na(s) + F2(g)  2NaF(s)
42
1. Reactions with Sodium
• Chlorine
 reacts violently to form sodium chloride
2Na(s) + Cl2(g)  2NaCl(s)
43
1. Reactions with Sodium
• Bromine
 burns steadily in bromine vapour
to form sodium bromide
2Na(s) + Br2(g)  2NaBr(s)
44
1. Reactions with Sodium
• Iodine
 burns steadily in iodine vapour to
form sodium iodide
2Na(s) + I2(g)  2NaI(s)
45
Na+(g) + X(g)
Vigor of reaction depends on
1
+(g) + X(g)
1. The
activation
energy
(endothermic)
B.E.
Na
2
Na+(g) +
1
2 X2(g)
2. The lattice energy (exothermic)
H
o
atm
 Activation energy
Hovap
Na(s) +
1
2 X2
Hof
46
Holattice
I.E.
NaX(s)
E.A.
Na+(g) + X(g)
1
2 B.E.
F is the most reactive
Na+(g) +
1
2 X2(g)
Hoatm
Na(s) +
Hof
47
Na+(g) + X(g)
F has an exceptionally
low B.E. & zero Hovap
Hovap
1
2 X2(g)
E.A.
Holattice
I.E.
NaX(s)
Na+(g) + X(g)
1
2 B.E.
E.A.
Na+(g) +
Na+(g) + X(g)
1
2 X2(g)
H
o
atm
The lattice enthalpy of NaF
is most negative
Hovap
Na(s) +
1
2 X2
Hof
48
Holattice
I.E.
NaX(s)
Na+(g) + X(g)
1
2 B.E.
Cl is more reactive than Br & I
E.A.
Na+(g) + X(g)
1
Na+(g) + 2 X2(g)
Cl has zero Hovap
Hoatm
Hovap
Na(s) +
1
2 X2(g)
Hof
49
Holattice
I.E.
NaX(s)
Na+(g) + X(g)
1
2 B.E.
E.A.
Na+(g) +
Na+(g) + X(g)
1
2 X2(g)
Lattice enthalpy :
Hoatm NaCl > NaBr > NaI
Hovap
Ho
lattice
Na(s) +
1
2 X2
Hof
50
I.E.
NaX(s)
Na+(g) + X(g)
1
2 B.E.
Br is more reactive than I
Na+(g) +
Na+(g) + X(g)
1
2 X2(g)
Hoatm
Hovap : Br2(l) < I2(s)
Hovap
Na(s) +
1
2 X2(s)/(l)
Hof
51
E.A.
Holattice
I.E.
NaX(s)
Na+(g) + X(g)
1
2 B.E.
E.A.
Na+(g) +
Na+(g) + X(g)
1
2 X2(g)
Lattice enthalpy :
Hoatm
NaBr > NaI
Hovap
Na(s) +
1
2 X2
Hof
52
Holattice
I.E.
NaX(s)
Q.2(a)
Variation: bond enthalpy decreases from Cl2 to I2
Reason :
The size of atoms and thus the bond
length between atoms increases down
the group.
The shared electron pair is getting
further away from the bonding nuclei.
 weaker bond and lower B.E.
F2 has an exceptionally small B.E. because the F atoms
are so small that the repulsive forces between lone
pairs on adjacent bonding atoms become significant.
53
Q.2(b)
The lattice enthalpy becomes less negative down the
group.
It is because the anionic radius, r- , increases down the
group.
H
54
o
lattice
1

r  r
2.1 Reactions with hydrogen
X2 + H2(g)  2HX(g)
F2 reacts explosively even in the dark at 200C
Cl2 reacts explosively in sunlight
Br2 reacts moderately on heating with a catalyst
I2 reacts slowly and reversibly even on heating
55
Q.3
Explain the extreme reactivity of fluorine in terms of
the bond enthalpies of F–F and H–F bonds.
F2 + H2(g)  2HF(g)
Fluorine has an exceptionally small F-F bond enthalpy.
Thus, the activation energy of its reaction with hydrogen is
also exceptionally small.
Hydrogen fluoride has the highest bond enthalpy among the
hydrogen halides.
Thus, the formation of HF from H2 and F2 is the most
exothermic.
The energy released from the reaction further speeds up the
reaction.
56
Chlorine removes hydrogen completely from
turpentine(C10H16)
C10H16(l) + 8Cl2(g)  10C(s) + 16HCl(g)
57
Q.4
The cotton wool bursts into flames and
the gas jar is filled with dark smoke (of
carbon) and white fumes (of HCl)
HCl gives dense white fumes with
ammonia.
58
2.2 Reactions with phosphorus
F2 + P  PF5
Cl2 + P  PCl3 + PCl5
Br2 + P  PBr3
I2 + P  PI3
F2 is the strongest oxidizing agent, it always
oxidizes other elements to their highest
possible oxidation states.
59
2.2 Reactions with phosphorus
F2 + P  PF5
Cl2 + P  PCl3 + PCl5
Br2 + P  PBr3
I2 + P  PI3
Br2 and I2 are NOT strong enough to oxidize
P to its highest possible oxidation state.
60
2.3 Reactions with xenon
Fluorine reacts directly with all non-metals except
nitrogen, helium, neon and argon.
It will even react with diamond and xenon on
heating.
C(diamond) + 2F2  CF4
Xe + F2  XeF2
Xe + 2F2  XeF4
Xe + 3F2  XeF6
61
2.3 Reactions with xenon
It is because
(a) Xenon can expand its octet by utilizing vacant,
low-lying d-orbitals.
62
By VB Theory,
5s
Xe 
5p
  
2s
F 
2p
  
To form two Xe-F bonds in XeF2, a 5p electron
in Xe has to be promoted to a 5d orbital.
5s
Xe* 
63
5d
5p
  

By VB Theory,
Xe
5d
5s
5p

  
To form four Xe-F bonds in XeF4, two 5p
electrons in Xe have to be promoted to two 5d
orbitals.
5d
5s
5p
Xe** 
64
 



By VB Theory,
Xe
5d
5s
5p

  
To form six Xe-F bonds in XeF6, three 5p
electrons in Xe have to be promoted to three 5d
orbitals.
5d
5s
5p
Xe*** 
65






2.3 Reactions with xenon
The gap between np and nd sub-shells  down
the group, thus,
the promotion of electrons from np sub-shell to
nd sub-shell becomes easier down the group.
Tendency to form bonds  down the group : Xe > Kr > Ar > Ne > He
66
Xe
5d
5s
5p

  
5s
Xe*** 
5d
5p






Also, the energy released by forming
more single bonds outweighs the energy
required for promoting 5p electrons to 5d
orbitals.
67
3
Reactions with other reducing agents
I2 is the weakest oxidizing agents among the
halogens.
68
3.1
All halogens(except I2) oxidize Fe2+ to Fe3+
Half reaction
Standard electrode
potential (V)
Cl2(aq) + 2e–  2Cl–(aq)
+1.36
Br2(aq) + 2e –  2Br–(aq)
+1.07
Fe3+(aq) + e–  Fe2+(aq)
+0.77
I2(aq) + 2e–  2I–(aq)
+0.54
X2(aq) + 2Fe2+(aq)  2X(aq) + 2Fe3+(aq)
o
0
( X = F, Cl, Br) Ecell
69
3.1
All halogens(except I2) oxidize Fe2+ to Fe3+
Half reaction
Standard electrode
potential (V)
Cl2(aq) + 2e–  2Cl–(aq)
+1.36
Br2(aq) + 2e –  2Br–(aq)
+1.07
Fe3+(aq) + e–  Fe2+(aq)
+0.77
I2(aq) + 2e–  2I–(aq)
+0.54
o
I2(aq) + 2Fe2+(aq)  No reaction Ecell
0
70
3.2
All halogens(except I2) oxidize S2O32 to
SO42
4X2(aq) + S2O32(aq) + 5H2O(l)  8X(aq) + 10H+(aq) + 2SO42(aq)
(X = F, Cl, Br)
S0
I2(aq) + 2S2O32(aq)  2I(aq) + S4O62(aq)
Used in iodometric titration
O
O-
O
0
+5 S
O
0
S
S
+5
O
+6
S
O
-
O
OO-
O
S
O
71
S +4
OO-
Determination of [Fe3+(aq)] by iodometric titration
(i)
(ii)
2I(aq) + 2Fe3+(aq)  I2(aq) + 2Fe2+(aq)
(excess) (unknown)
I2(aq) + 2S2O32(aq)  2I(aq) + S4O62(aq)
(standard solution)
nFe3 : nS O 2  1 : 1
2 3
Using starch as indicator
72
4
Displacement reactions
Cl2(aq) + 2Br(aq)  2Cl(aq) + Br2(aq)
Cl2(aq) + 2I(aq)  2Cl(aq) + I2(aq)
Br2(aq) + 2I(aq)  2Br(aq) + I2(aq)
More reactive
Less reactive
I2(aq) + I(aq)  I3(aq)
(yellow)
73
(brown)
4
Displacement reactions
Cl2(aq) + 2I(aq)  2Cl(aq) + I2(aq)
Br2(aq) + 2I(aq)  2Br(aq) + I2(aq)
I2(aq) + I(aq)  I3(aq)
(yellow)
(brown)
What would be observed if an excess of
Cl2(aq) or Br2(aq) is added to I(aq)?
74
The solution turns cloudy and a black
solid settles at the bottom
Reactions of halide ions with halogens
Aqueous
solution
75
Halogen added
F2
F–
No reaction
Cl–
A pale yellow
solution is
formed (Cl2
is formed)
Cl2
Br2
I2
No reaction
No
reaction
No
reaction
No
reaction
No
reaction
No reaction
Reactions of halide ions with halogens
Aqueous
solution
Halogen added
F2
A yellow
solution
Br–
is formed
(Br2 is
formed)
I–
76
Cl2
A yellow
solution
is formed
(Br2 is formed)
Br2
I2
No reaction
No
reaction
A yellowish
brown
A yellowish
brown
A yellowish
brown
solution is
formed
solution is
formed
solution is
formed
(I3 is formed)
(I3 is formed)
(I3 is formed)
No
reaction
Q.5
Shake hexane or 1,1,1-trichloroethane with
the two solutions respectively.
The one that turns the organic layer violet is
I3(aq).
The one that turns the organic layer orange or
brown is Br2(aq).
77
If hexane is used, the upper layer will be the organic layer
Br2(aq)
I3(aq)
1,1,1-trichloroethane
Br2
78
I2
5.
Disproportionation
Disproportionation is a chemical change in
which oxidation and reduction of the same
species (which may be a molecule, atom or
ion) take place at the same time.
79
A. Reactions with Water
HOCl : chloric(I) acid or hypochlorous acid
Chlorine water
80
 a mixture of hydrochloric acid and
chloric(I) acid
A. Reactions with Water
• Chlorate(I) ion, OCl is also known as
hypochlorite ion
 unstable
 decomposes when exposed to
sunlight or high temperatures to give
chloride ions and oxygen
2OCl–(aq)  2Cl–(aq) + O2(g)
81
A. Reactions with Water
• Chlorate(I) ion
 bleaches by oxidation
Cl2(aq) + H2O(l)
2H+(aq) + Cl–(aq) + OCl–(aq)
OCl–(aq) + dye  Cl–(aq) + (dye + O)
coloured
82
colourless
A. Reactions with Water
• Bromine
 only slightly soluble in water
 mainly exists as molecules in
saturated bromine water
83
A. Reactions with Water
• When the solution is diluted
 hydrolysis takes place
 hydrobromic acid and bromic(I) acid
(hydrobromous acid) are formed
Br2(l) + H2O(l)
84
HBr(aq) + HOBr(aq)
A. Reactions with Water
• Bromate(I) ion, OBr
 also unstable
 bleaches dyes by oxidation
OBr–(aq) + dye
coloured
 Br–(aq) + (dye + O)
colourless
85
A. Reactions with Water
• Iodine
 does not react with water
 only slightly soluble in water
86
A. Reactions with Water
• Fluorine
reacts vigorously with water to form
hydrogen fluoride and oxygen
0
1
2F2(g) + 2H2O(l)  4HF(aq) + O2(g)
Being the strongest oxidizing agent,
F2 undergoes reduction rather than
disproportionation with water.
87
A. Reactions with Water
Chlorine reacts similarly at high temperature
or when exposed to light
2Cl2(aq) + 2H2O(l)  2HCl(aq) + 2HOCl(aq)
2HOCl(aq)
Heat or light
2HCl(aq) + O2(g)
Overall :
2Cl2(aq) + 2H2O(l)
88
Heat or light
4HCl(aq) + O2(g)
B. Reactions with Alkalis
• All halogens react with aqueous alkalis
• All halogens (except F2) undergoes
disproportionation with alkalis
• In general,
Reactivity decreases down the group
89
B. Reactions with Alkalis
The products formed depend on
1. Temperature
2. The type of halogen reacted
3. The concentration of alkali used
90
B. Reactions with Alkalis
Effect of temperature
(a) At lower temperatures,
T1
0
X2(aq) + 2OH(aq)
91
1
+1
XO(aq) + X(aq) + H2O(l)
X2
Cl2
Br2
I2
T1 / C
20
0
<0
B. Reactions with Alkalis
Effect of temperature
(a) At higher temperatures,
+1
3XO(aq)
92
T2
+5
1
XO3(aq) + 2X(aq)
XO
ClO
BrO
IO
T2 / C
70
20
0
B. Reactions with Alkalis
(1)
X2(aq) +
(2)
3XO(aq)
T1
2OH(aq)
T2
XO(aq) + X(aq) + H2O(l)
XO3(aq) + 2X(aq)
Overall reaction : 3(1) + (2)
3X2(aq) +
93
6OH(aq)
T2
XO3(aq) + 5X(aq) + 3H2O(l)
B. Reactions with Alkalis
3XO(aq)
T2
XO3(aq) + 2X(aq)
XO
ClO
BrO
IO
T2 / C
70
20
<0
On moving down the group,
94
1. stability of XO decreases
ClO > BrO > IO
2. stability of XO3 increases
ClO3 < BrO3 < IO3
B. Reactions with Alkalis
3X2(aq) + 6OH(aq)
XO3(aq) + 5X(aq) + 3H2O(l)
At lower pH (when acid is added),
the equilibrium position shifts to the left and the
reversed process predominates.
XO3(aq) + 5X(aq) + 6H+(aq)
3X2(aq) + 3H2O(l)
This reaction (when X=I) is often used to prepare
standard iodine solution for iodometric titrations
95
B. Reactions with Alkalis
• Dissolving a known quantity of KIO3(s)
in excess KI(aq) and dilute H2SO4
generates a known amount of I2(aq)
KIO3(aq) + 5KI(aq) + 6H+(aq)
 3I2(aq) + 3H2O(l) + 6K+(aq)
• The iodine produced can be used to
standardize thiosulphate solution
3I2(aq) + 6S2O32(aq)
 6I(aq) + 3S4O62(aq)
96
B. Reactions with Alkalis
•
This known amount of iodine generated can also
be used to oxidize reducing agents (of unknown
concentrations) such as SO32(aq) and ascorbic
acid (vitamin C)
•
The excess iodine can be determined by
back titration with sodium thiosulphate
solution
I2(aq) + 2S2O32–(aq)
 2I–(aq) + S4O62–(aq)
97
B. Reactions with Alkalis
Effect of concentration of alkali
(a) At higher concentrations, XO3(aq) is the
major product.
(b)At lower concentrations, XO(aq) is the
major product.
98
B. Reactions with Alkalis
In general,
Halogens react with cold, dilute alkali to give
halate(I) ions, halide ions and water
X2(aq) + 2OH  XO(aq) + X(aq) + H2O(l)
Halogens react with hot, concentrated alkali to give
halate(V) ions, halide ions and water.
3X2(aq) + 6OH  XO3(aq) + 5X(aq) + 3H2O(l)
99
B. Reactions with Alkalis
0
2
2F2 + 2OH(aq)
20C
+2 1
1
OF2(aq) + 2F(aq) + H2O(l)
very dilute
0
2
2F2 + 4OH(aq)
70C
0
1
O2(aq) + 4F(aq) + 2H2O(l)
concentrated
Being the strongest oxidizing agent, F2 undergoes
reduction rather than disproportionation with alkalis.
100
Variation in chemical properties of halides
A Comparative study
1. Reactions with conc. sulphuric acid
2. Reactions with conc. phosphoric acid
3. Reactions with silver ion
101
Reactions with Concentrated
Sulphuric(VI) Acid
•
Concentrated sulphuric acid
 non-volatile (b.p. ~330C)
 oxidizing
102
Fluoride and chloride : warm
KF(s) + H2SO4(l)  KHSO4(s) + HF(g)
warm
KCl(s) + H2SO4(l)  KHSO4(s) + HCl(g)
non-volatile
volatile
Warming is required to speed up the reaction and to
drive out the volatile acids
103
warm
KF(s) + H2SO4(l)  KHSO4(s) + HF(g)
warm
KCl(s) + H2SO4(l)  KHSO4(s) + HCl(g)
acid salt
Acid salt rather than normal salt is formed because
HSO4 is a relatively weak acid
A convenient way to prepare HCl in the laboratory
104
warm
KF(s) + H2SO4(l)  KHSO4(s) + HF(g)
warm
KCl(s) + H2SO4(l)  KHSO4(s) + HCl(g)
Observation : White fumes are produced
Confirmatory test : Dense white fumes appear with NH3(aq)
105
Bromide: warm
KBr(s) + H2SO4(l)  KHSO4(s) + HBr(g)
reduction
-1
+6
warm
+4
0
2HBr(g) + H2SO4(l)  SO2(g) + Br2(g) + 2H2O(l)
oxidation
106
Bromide: warm
(1) KBr(s) + H2SO4(l)  KHSO4(s) + HBr(g)
warm
(2) 2HBr(g) + H2SO4(l)  SO2(g) + Br2(g) + 2H2O(l)
Overall reaction : 2(1) + (2)
2KBr(s) + 3H2SO4(l)
warm
 2KHSO4(s) + SO2(g) + Br2(g) + 2H2O(l)
Not suitable for preparing HBr
107
2KBr(s) + 3H2SO4(l)  2KHSO4(s) + SO2(g) + Br2(g) + 2H2O(l)
Halide
Br–
Observation
Product
• White fumes
are formed
HBr
• Dense white fumes are
formed with aqueous
ammonia
• A pungent
smell is
SO2
• It turns orange
dichromate solution
detected
• A brown gas is
evolved on
warming
108
Confirmatory Test
green
Br2
• A brown colour is
observed when adding
hexane
iodide: warm
KI(s) + H2SO4(l)  KHSO4(s) + HI(g)
-1
+6
warm
-1
+6
warm
+4
0
2HI(g) + H2SO4(l)  SO2(g) + I2(g) + 2H2O(l)
-2
0
8HI(g) + H2SO4(l)  H2S(g) + 4I2(g) + 2H2O(l)
HI is strong enough to reduce sulphur to its
lowest possible oxidation state
109
warm
KI(s) + H2SO4(l)  KHSO4(s) + HI(g) (1)
warm
2HI(g) + H2SO4(l)  SO2(g) + I2(s) + 2H2O(l) (2)
warm
8HI(g) + H2SO4(l)  H2S(g) + 4I2(s) + 2H2O(l) (3)
Overall reaction = 10(1) + (2) + (3)
10KI(s) + 12H2SO4(l)
 10KHSO4(s) + SO2(g) + H2S(g) + 5I2(s) + 4H2O(l)
No suitable for preparing HI
110
10KI(s) + 12H2SO4(l)
 10KHSO4(s) + SO2(g) + H2S(s) + 5I2(s) + 4H2O(l)
Observation : A bad egg smell is detected
Confirmatory test : It turns lead(II) ethanoate paper black
(CH3COO)2Pb + H2S  PbS(s) + 2CH3COOH
111
10KI(s) + 12H2SO4(l)
 10KHSO4(s) + SO2(g) + H2S(s) + 5I2(s) + 4H2O(l)
Observation : Violet fumes are formed and
condense when cooled to give a black solid
Confirmatory test : A violet colour is observed when added to
hexane
112
Conclusion : -
Increases down the group
Reducing power : HI > HBr > HCl > HF
113
Interpretation:Consider the reaction,
2H–X + H2SO4  X–X + SO2 + 2H2O
The feasibility of the reaction depends on
1. the strength of H–X bond to be broken
the stronger the bond, the less feasible is the rx
2. the strength of X–X bond to be formed
the stronger the bond, the more feasible is the rx
114
2H–X + H2SO4  X–X + SO2 + 2H2O
The feasibility of the reaction depends on
1. the strength of H–X bond
the stronger the bond, the less feasible is the rx
2. the strength of X–X bond
the stronger the bond, the more feasible is the rx
The reaction with HF is least feasible because
1. H-F bond is the strongest
2. F-F bond is exceptionally weak due to repulsion
between lone pairs of bonding atoms.
115
H-X
B.E.(kJ mol1)
X-X
B.E. (kJ mol1
H-Cl
432
Cl-Cl
244
H-Br
366
Br-Br
192
H-I
298
I-I
152
On moving down the group,
both H–X bonds and X–X bonds become weaker
116
2H–X + H2SO4  X–X + SO2 + 2H2O
The strength of H-X bond is more important
Since two H-X bonds have to be broken for each
X-X bond formed.
Reactivity : H-Cl < H-Br < H-I
117
Reactions with Phosphoric Acid
warm
NaCl(s) + H3PO4(l)  NaH2PO4(s) + HCl(g)
warm
NaBr(s) + H3PO4(l) 
NaH2PO4(s) + HBr(g)
warm
NaI(s) + H3PO4(l)  NaH2PO4(s) + HI(g)
non-volatile
volatile
H3PO4(l) + HX(g)  no reaction
less oxidizing
Suitable for preparing HX from solid halids
118
warm
NaCl(s) + H3PO4(l)  NaH2PO4(s) + HCl(g)
warm
NaBr(s) + H3PO4(l) 
NaH2PO4(s) + HBr(g)
warm
NaI(s) + H3PO4(l)  NaH2PO4(s) + HI(g)
Halide
ion
Cl–
Br–
I–
119
Observation
White fumes are
formed on warming
Confirmatory test of
Product
the product
HCl
HBr
HI
Dense white fumes are
formed with aqueous
ammonia
Reactions with Silver Ions
•
Aqueous solutions of chlorides, bromides and
iodides
 give precipitates when reacting
with acidified silver nitrate solution
120
Reactions with Silver Ions
Ag+(aq) + Cl–(aq)  AgCl(s)
white ppt
Ag+(aq) + Br–(aq)  AgBr(s)
pale yellow ppt
Ag+(aq) + I–(aq)
121
 AgI(s)
yellow ppt
AgCl(s)
AgBr(s)
AgI(s)
Colour intensity  down the group
122
Reactions with Silver Ions
Silver nitrate solution should be acidified
with nitric acid
(a)
to remove interfering ions like
SO32 or CO32
They may form white ppt with Ag+
123
Reactions with Silver Ions
2H+(aq) + SO32–(aq)  SO2(g) + H2O(l)
2H+(aq) + CO32–(aq)  CO2(g) + H2O(l)
124
Silver nitrate solution should be acidified
with nitric acid
(b) to avoid the formation of black ppt
of Ag2O in alkaline solution.
2Ag+(aq) + 2OH(aq)  Ag2O(s) + H2O(l)
125
The solubility(in water) of AgX  down the group
soluble
AgF
insoluble
>>
Ksp/mol2 dm6
126
AgCl
1.61010
>
AgBr
7.71013
>
AgI
1.51016
Q.7
On moving down the group,
the size of the halide anions 
 The electron cloud of the anions becomes
more easily polarized by Ag+
 The halides become more covalent and
less ionic
 The halides become less soluble in polar
solvents like water
127
Reactions with Silver Ions
The reaction can be used as a test
to show the presence of halide ions.
Different halides give ppt with
different colours.
Sometimes ambiguous.
Confirmatory tests are needed.
128
Two confirmatory tests for halides
1.Adding NH3(aq) to the AgX ppt
2.Exposing AgX ppt to sunlight
129
AgX(s) dissolve in NH3(aq) due to the formation
of soluble complex ions.
AgCl(s) + 2NH3(aq)  [Ag(NH3)2]+(aq) + Cl(aq)
AgBr(s) + 2NH3(aq)  [Ag(NH3)2]+(aq) + Br(aq)
AgI(s) + 2NH3(aq)  No reaction
Solubility in NH3(aq)  down the group
130
•
When exposed to sunlight
 silver chloride turns grey
light
2AgCl(s)  2Ag(s) + Cl2(g)
 silver bromide turns yellowish grey
light
2AgBr(s)  2Ag(s) + Br2(l)
 silver iodide remains yellow
light
2AgI(s)  No reaction
131
Action of acidified silver nitrate solution on halides
Confirmatory test of the product
Action of
acidified
Effect of
Ion
Effect of adding
AgNO3 solution
exposure
aqueous ammonia
on halides
to sunlight
Cl–
Br–
I–
132
A white ppt is
formed
The white ppt
dissolves
The solution
turns grey
A pale yellow
ppt is formed
The pale yellow
ppt slightly
dissolves
The solution
turns
yellowish grey
A yellow ppt is
formed
The yellow ppt
does not dissolve
The solution
remains yellow
Anomalous Behaviour of Hydrogen
Fluoride
1. Hydrogen fluoride has abnormally high
boiling point and melting point among
the hydrogen halides
133
HX
HF
HCl
HBr
HI
b.p./C
19.5
85
66.4
35
•
Molecules of all other hydrogen halides
 held together by weak van der Waal’s
forces only
Formation of the extensive intermolecular hydrogen
bonds among hydrogen fluoride molecules
134
2.
Acidic Properties of Hydrogen Halides
•
The acid strength of hydrogen halides
decreases in the order:
HI > HBr > HCl >> HF
135
Acid dissociation constants of hydrogen halides and their
degrees of dissociation in 0.1 M solutions
Hydrogen
halide
Acid dissociation
constant,
Ka (mol dm–3)
Degree of
dissociation in 0.1 M
solution (%)
Acid
strength
HF
5.6 × 10–4
8.5
Low
HCl
1 × 107
92
Strong
HBr
1 × 109
93
Strong
HI
1 × 1011
95
Very strong
136
HF(l) + H2O(l)
H3O+(aq) + F–(aq)
Ka = 5.6 × 10–4 mol dm–3
In dilute (e.g. 0.1M) solution,
HF is the weakest acid among all the
hydrohalic acids
137
H-bond
H
H
O
H
+
F
O
H
F
H
H
or
H3O + F
H3OF
Very stable ion pair
Freedom of F & H3O+ greatly  (a drop in
entropy of the system) due to H-bond
formation
Effective concentration of F & H3O+ greatly 
Thus, Ka  & pH 
138
In concentrated solution,
HF is the strongest acid among all the
hydrohalic acids
139
Strength of H-bond:H
F
H
F
>
O
H
F
H
2.
HF is in excess in concentrated solution
F ions combine with excess HF
rather than with H3O+
 free H3O+  & pH 
H3OF(aq) + HF(aq)  H3O+(aq) + HF2(aq)
excess
140
For other HX acids,
acidity  as concentration 
It is due to the significant interaction
between X and H3O+ at high concentrations
 the effective concentration of H3O+ 
For HF, interaction between F and H3O+ is
significant even at low concentrations due
to the smaller size of F.
141
3.
Pure, anhydrous liquid HF is ionic due to
the formation of HF2 and H2F+ ions
Self ionization : 2HF(l)
HF(l) + F(l)
H2F+(l) + F(l)
HF2(l)
Overall : -
3HF(l)
142
[H2F]+[HF2](l)
F
H
F

F
H
F
Stabilized by resonance

Two identical H – F bonds
143
KF(s) + HF(l)
•
heat
KHF2(s)
Heating the solid potassium hydrogen
difluoride
 reverses the reaction
 a convenient way to obtain
anhydrous hydrogen fluoride
144
Uses of fluorine and its compounds
Sodium hexafluorosilicate, Na2SiF6, is used in
water fluoridation.
F, being isoelectronic to OH, can replace the
OH in the tooth enamel, making it less soluble
in acidic solutions.
145
Uses of fluorine and its compounds
Molten cryolite, Na3AlF6
Lowers the temperature (2517C  1000 C)
needed for extracting Al from Al2O3 by
electrolysis.
146
Uses of fluorine and its compounds
Convert U to UF6
Separate 235UF6 from 238UF6 by diffusion for
use in nuclear reactors.
The heavier 238UF6 diffuses a bit slower,
making the separation possible.
147
Uses of fluorine and its compounds
Conc. HF(aq) is used in etching glass
(e.g. making scales/graduation marks on
glassware)
CaSiO3(s) + 6HF(aq)
(Glass)
148
 CaF2(aq) + SiF4(aq) + 3H2O(l)
Uses of fluorine and its compounds
•
The glass object to be etched
 coated with wax or a similar acidproof material
 cutting through the wax layer to
expose the glass
 apply hydrofluoric acid
149
Uses of fluorine and its compounds
A glass is etched by
hydrofluoric acid
150
Uses of fluorine and its compounds
Making fluorocarbon compounds
Used as refrigerants, aerosol propellants,
anaesthetics and
fire-fighting agents(BTM, BCF)
PTFE (teflon) used in electrical insulation,
coating on surface of non-stick saucepans, etc.
151
Uses of fluorine and its compounds
Hydrazine/fluorine mixtures are excellent
rocket fuels
N2H4(g) + 2F2(g)  N2(g) + 4HF(g)
H = -1166 kJ mol1 (extremely exothermic)
Due to the strong NN and H-F bonds
152
Uses of fluorine and its compounds
Extraction of fluorine
Electrolyte : KF(s) dissolved in pure HF(l)
Anode : graphite
Cathode : steel
153
Q.8(a)
Anode : 2HF2  2HF + F2 + 2e
Cathode : 2H2F+ + 2e  2HF + H2
Overall :
2HF2 + 2H2F+  4HF + F2 + H2
154
8.(b) Overall :
2HF2 + 2H2F+  4HF + F2 + H2
6HF  4HF + F2 + H2
2HF  F2 + H2
KF is added to increase the conductivity of
the electrolyte.
KHF2 > HF or [H2F][HF2]
155
Q.8(c)
OH- (from H2O) rather than HF2 is oxidized
at the anode
Also, F2 reacts vigorously with water.
2F2(g) + 2H2O(l)  4HF(aq) + O2(g)
vigorous reaction
156
8.(d)
At high temperatures, fluorine produced can
react vigorously with the electrodes, air, etc.
157
Uses of Chlorine and its compounds
Polyvinyl chloride, PVC
making electrical insulation, bottles, floor tiles,
table cloth, shower curtain, etc.
158
CH2=CH2 + Cl2  CH2Cl – CH2Cl
CH2Cl – CH2Cl
heat
CH2=CHCl + HCl
Cl
n(CH2=CHCl)
159

C
H2
C
H
n
Making chlorine bleach
Cl2(g) + 2NaOH(aq)  NaCl(aq) + NaOCl + H2O(l)
Disinfectant in sterilizing water and sewage
treatment.
Extraction of bromine from sea water
Cl2(g) + 2Br(aq)  2Cl(aq) + Br2(aq)
160
Uses of Bromine and its compounds
Manufacture of 1,2-dibromoethane to remove
Pb from petrol engine
Pb(C2H5)4, TEL : anti-knock agent added to
petrol engine to prevent premature ignition.
TEL decomposes to give Pb that may cause
damage to the engine
Air pollutant
CH2Br-CH2Br + Pb(C2H5)4  PbBr2
161
volatile and emitted to air easily
AgBr is used in black-and-white photography
exposure to
light
coated on film
2AgBr(s)
2Ag(s) + Br2(l)
black
The excess AgBr(s) is removed as soluble
complex ion.
AgBr(s) + 2S2O32(aq)  [Ag(S2O3)2]3(aq) + Br(aq)
hypo
162
Uses of Iodine and its compounds
Making iodine tincture (antiseptic)
I2 in alcohol or KI(aq)
Radioactive iodine-131 as tracer in medical
diagnosis
Iodide is used to make iodized table salt for
preventing development of goitre.
163
Laboratory preparation of halogens(except F2)
conc. H2SO4
MnO2 +
NaCl
164
-1
+4
2NaCl + MnO2 + 2H2SO4
+2
0
 Na2SO4 + MnSO4 + 2H2O + Cl2
NaCl + H2SO4  HCl + NaHSO4
conc. H2SO4
Free from HCl and H2O
MnO2 +
NaCl
165
To remove HCl
To dry Cl2
Laboratory preparation of halogens(except F2)
conc. HCl
MnO2
166
Laboratory preparation of halogens(except F2)
conc. HCl
MnO4
167
The END
168