The Transition Metals: Properties and Applications

2
Why we study the Transition
Metals
 Transition metals are found in nature
◦ Rocks and minerals contain transition metals
◦ The color of many gemstones is due to the
presence of transition metal ions
 Rubies are red due to Cr
 Sapphires are blue due to presence of Fe and Ti
◦ Many biomolecules contain transition metals that
are involved in the functions of these biomolecules
 Vitamin B12 contains Co
 Hemoglobin, myoglobin, and cytochrome C contain Fe

3
Why we study the Transition
Metals
 Transition metals and their compounds
have many useful applications
◦ Fe is used to make steel and stainless steel
◦ Ti is used to make lightweight alloys
◦ Transition metal compounds are used as
pigments
 TiO2 = white
 PbCrO4 = yellow
 Fe4[Fe(CN)6]3 (prussian blue)= blue
◦ Transition metal compounds are used in
many industrial processes

4
Introduction
• d-block elements
◦ The elements of periodic table belonging
to group 3 to 12 are known as d-Block
elements. because in these elements last
electron enters in d sub shell or d orbital
◦ locate between the s-block and p-block
◦ occur in the fourth and subsequent
periods of the Periodic Table

6
period 4
period 5
period 6
period 7
d-block elements

7
Transition elements are elements that
contain an incomplete d sub-shell (i.e. d1
to d9) in at least one of the oxidation states
of their compounds.
3d0
3d10
Introduction

8
How are d - Block Elements &
Transition elements different?
Not all d block elements are transition
elements but all transition elements are
d-block elements
Not all d block elements are transition
elements because d block elements like
Zinc have full d10 configuration in their
ground state as well as in their common
oxidation state which is not according to
definition of transition elements.

9
Introduction
Sc and Zn are not transition elements
because
They form compounds with only one
oxidation state in which the d sub-shell are
NOT incomplete.
Sc  Sc3+ 3d0 Zn  Zn2+ 3d10

10
Introduction
Cu
Cu+ 3d10
not transitional
Cu2+ 3d9
transitional

1. Which of the d-block elements may not be
regarded as the transition elements?
2. Why Zn, Cd and Hg are not considered as
transition elements.
3. Copper atom has completely filled d orbital
(3d10) in its ground state, yet it is transition
element. Why
4. Silver atom has completely filled d orbital
(4d10) in its ground state, yet it is transition
element. Why
5. Why the very name ‘transition’ given to the
elements of d-block .

1. Zn, Cd and Hg
2. Because they do not have vacant d-orbitals
neither in the atomic state nor in any stable
oxidation state.
3. Copper (Z = 29) can exhibit +2 oxidation state
wherein it will have incompletely filled d-orbitals
(3d), hence a transition element.
4. Silver (Z = 47) can exhibit +2 oxidation state
wherein it will have incompletely filled d-orbitals
(4d), hence a transition element.
5. The very name ‘transition’ given to the elements
of d-block is only because of their position
between s– and p– block elements.
Answers

14
The first transition series
the first horizontal row of the d-block elements

15
Characteristics of transition elements
(d-block metals vs s-block metals)
1. Physical properties vary slightly with atomic
number across the series (cf. s-block and
p-block elements)
2. Higher m.p./b.p./density/hardness than
s-block elements of the same periods.
3. Variable oxidation states
(cf. fixed oxidation states of s-block metals)

16
Characteristics of transition elements
4. Formation of coloured compounds/ions
(cf. colourless ions of s-block elements)
5. Formation of complexes
6. Catalytic properties

17
The building up of electronic configurations
of elements follow:
 Aufbau principle
 Hund’s rule
 Pauli exclusion principle
Electronic Configurations

18
• 3d and 4s sub-shells are very close to
each other in energy.
• Relative energy of electrons in sub-
shells depends on the effective nuclear
charge they experience.
• Electrons enter 4s sub-shell first
• Electrons leave 4s sub-shell first

19
Cu Cu2+
Relative energy levels of orbitals
in atom and in ion

20
• Valence electrons in the inner 3d orbitals
• Examples:
 The electronic configuration of
scandium: 1s22s22p63s23p63d14s2
 The electronic configuration of zinc:
1s22s22p63s23p63d104s2

21
Element Atomic number Electronic configuration
Scandium
Titanium
Vanadium
Chromium
Manganese
Iron
Cobalt
Nickel
Copper
Zinc
21
22
23
24
25
26
27
28
29
30
[Ar] 3d 14s2
[Ar] 3d 24s2
[Ar] 3d 34s2
[Ar] 3d 54s1
[Ar] 3d 54s2
[Ar] 3d 64s2
[Ar] 3d 74s2
[Ar] 3d 84s2
[Ar] 3d 104s1
[Ar] 3d 104s2
Electronic configurations of the first series of the
d-block elements

22
• A half-filled or fully-filled d sub-shell
has extra stability

23
d -Block Elements as Metals
Physical properties of d-Block elements :
 good conductors of heat and electricity
 hard and strong
 malleable and ductile
• d-Block elements are typical metals

24
• Physical properties of d-Block elements:
• Exceptions : Mercury
 low melting point
 liquid at room temperature and
pressure
 lustrous
 high melting points and boiling points

25
• d-block elements
 extremely useful as construction
materials
 strong and unreactive

26
 used for construction and making
machinery nowadays
 abundant
 easy to extract
• Iron
cheap

27
• Iron
 corrodes easily
 often combined with other
elements to form steel
 harder and more resistant to
corrosion

28
• Titanium
 used to make aircraft and space
shuttles
 expensive
Corrosion resistant, light, strong and
withstand large temperature changes

29
• Manganese
confers hardness & wearing resistance to
its alloys
e.g. duralumin : alloy of Al with Mn/Mg/Cu
• Chromium
 confers inertness to stainless steel

30
• The similar atomic radii of the
transition metals facilitate the
formation of substitutional alloys
 the atoms of one element to
replace those of another element
 modify their solid structures and
physical properties

31
Atomic Radii and Ionic Radii
• Two features can be observed:
1. The d-block elements have smaller
atomic radii than the s-block elements
2. The atomic radii of the d-block
elements do not show much variation
across the series

32
Variation in atomic radius
of the first 36 elements

35
(i) Nuclear charge 
(ii) Shielding effect (repulsion between e-) 
(i) > (ii)
(i)  (ii)
(ii) > (i)
On moving across the Period,

36
• At the beginning of the series
 atomic number 
 effective nuclear charge 
 the electron clouds are pulled
closer to the nucleus
 atomic size 

37
• In the middle of the series
 the effective nuclear charge
experienced by 4s electrons increases
very slowly
 only a slow decrease in atomic radius
in this region
 more electrons enter the inner
3d sub-shell
 The inner 3d electrons shield the
outer 4s electrons effectively

38
• At the end of the series
 the screening and repulsive effects
of the electrons in the 3d sub-
shell become even stronger
 Atomic size 

39
• Many of the differences in physical and
chemical properties between the d-block
and s-block elements
 explained in terms of their differences
in electronic configurations and
atomic radii
Comparison of Some Physical and
Chemical Properties between the
d-Block and s-Block Elements

40
1. Density
Densities (in g cm–3) of the s-block elements and
the first series of the d-block elements at 20C

41
• d-block > s-block
 the atoms of the d-block elements
1. are generally smaller in size
2. are more closely packed
(fcc/hcp vs bcc in group 1)
3. have higher relative atomic masses
1. Density

42
• The densities
 generally increase across the first
series of the d-block elements
 1. general decrease in atomic
radius across the series
2. general increase in atomic mass
across the series
1. Density

43
2. Ionization Enthalpy
Element
Ionization enthalpy (kJ mol–1)
1st 2nd 3rd 4th
K
Ca
418
590
3 070
1 150
4 600
4 940
5 860
6 480
Sc
Ti
V
Cr
632
661
648
653
1 240
1 310
1 370
1 590
2 390
2 720
2 870
2 990
7 110
4 170
4 600
4 770
K  Ca (sharp ) ; Ca  Sc (slight )

44
Element
Ionization enthalpy (kJ mol–1)
1st 2nd 3rd 4th
Cr
Mn
Fe
Co
Ni
Cu
Zn
653
716
762
757
736
745
908
1 590
1 510
1 560
1 640
1 750
1 960
1 730
2 990
3 250
2 960
3 230
3 390
3 550
3 828
4 770
5 190
5 400
5 100
5 400
5 690
5 980
Sc  Cu (slight ) ; Cu  Zn (sharp )

45
• The first ionization enthalpies of the
d-block elements
 greater than those of the s-block
elements in the same period of the
Periodic Table
 1. The atoms of the d-block
elements are smaller in size
2. greater effective nuclear charges

46
Sharp  across periods 1, 2 and 3
Slight  across the transition series

47
• Going across the first transition series
 the nuclear charge of the elements
increases
 additional electrons are added to
the ‘inner’ 3d sub-shell

48
• The screening effect of the additional
3d electrons is significant
• The effective nuclear charge experienced
by the 4s electrons increases very slightly
across the series
• For 2nd, 3rd, 4th… ionization enthalpies,
slight and gradual  across the series
are observed.

49
Electron has to be removed from
completely filled 3p subshell
3d5
3d5
3d5
3d10
d10/s2Cr+
Mn2+
Fe3+

50
• The first few successive ionization
enthalpies for the d-block elements
 do not show dramatic changes
 4s and 3d energy levels are close to
each other

51
3. Melting Points and Hardness
1541 1668 1910 1907 1246 1538 1495 1455 1084 419
d-block >> s-block
 1. both 4s and 3d e- are involved in the
formation of metal bonds
2. d-block atoms are smaller

52
K has an exceptionally small m.p. because it has an
more open b.c.c. structure.
1541 1668 1910 1907 1246 1538 1495 1455 1084 419

53
 Unpaired electrons are relatively
more involved in the sea of electrons
Sc Ti V Cr Mn Fe Co Ni Cu Zn
1541 1668 1910 1907 1246 1538 1495 1455 1084 419

54
 
3d 4s
Sc
  Ti
   V
1. m.p.  from Sc to V due to the  of
unpaired d-electrons (from d1 to d3)
1541 1668 1910 1907 1246 1538 1495 1455 1084 419

55
2.m.p.  from Fe to Zn due to the 
of unpaired d-electrons (from 4 to 0)
1541 1668 1910 1907 1246 1538 1495 1455 1084 419
     
3d 4s
Fe
     Co
     Ni

56
1541 1668 1910 1907 1246 1538 1495 1455 1084 419
3. Cr has the highest no. of unpaired
electrons but its m.p. is lower than V.
     
3d 4s
Cr
It is because the electrons in the
half-filled d-subshell are relatively
less involved in the sea of electrons.

57
1541 1668 1910 1907 1246 1538 1495 1455 1084 419
4. Mn has an exceptionally low m.p.
because it has the very open cubic
structure.
Why is Hg a liquid at room conditions ?
All 5d and 6s electrons are paired up
and the size of the atoms is much
larger than that of Zn.

58
• The metallic bonds of the d-block
elements are stronger than those of the
s-block elements
 much harder than the s-block
elements
• The hardness of a metal depends on
 the strength of the metallic bonds

59
Mohs scale : - A measure of hardness
Talc Diamond
0 10
K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn
0.5 1.5 3.0 4.5 6.1 9.0 5.0 4.5 -- -- 2.8 2.5

60
• In general, the s-block elements
 react vigorously with water to form
metal hydroxides and hydrogen
4. Reaction with Water
• The d-block elements
 react very slowly with cold water
 react with steam to give metal oxides
and hydrogen

61
4. Reaction with Water
2K(s) + 2H2O(l)  2KOH(aq) + H2(g)
2Na(s) + 2H2O(l)  2NaOH(aq) + H2(g)
Ca(s) + 2H2O(l)  Ca(OH)2(aq) + H2(g)
Zn(s) + H2O(g)  ZnO(s) + H2(g)
3Fe(s) + 4H2O(g)  Fe3O4(s) + 4H2(g)

62
d-block compounds vs s-block compounds
A Summary : -
Ions of d-block metals have higher charge density
 more polarizing
 1. more covalent in nature
2. less soluble in water
3. less basic (more acidic)
Basicity : Fe(OH)3 < Fe(OH)2 << NaOH
Charge density : Fe3+ > Fe2+ > Na+

63
d-block compounds vs s-block compounds
A Summary : -
4. less thermally stable e.g. CuCO3 << Na2CO3
5. tend to exist as hydrated salts
e.g. CuSO4.5H2O, CoCl2.2H2O
6. hydrated ions undergo hydrolysis more easily
e.g. [Fe(H2O)6]3+(aq) + H2O  [Fe(OH)(H2O)5]2+(aq) + H3O+
acidic

64
• One of the most striking properties
 variable oxidation states
Variable Oxidation States
• The 3d and 4s electrons are
 in similar energy levels
 available for bonding

65
• Elements of the first transition series
 form ions of roughly the same
stability by losing different
numbers of the 3d and 4s electrons
Variable Oxidation States

66
Oxidation
states
Oxides / Chloride
+1
Cu2O
Cu2Cl2
+2
TiO VO CrO MnO FeO CoO NiO CuO ZnO
TiCl2 VCl2 CrCl2 MnCl2 FeCl2 CoCl2 NiCl2 CuCl2 ZnCl2
+3
Sc2O3 Ti2O3 V2O3 Cr2O3 Mn2O3 Fe2O3 Ni2O3 • xH2O
ScCl3 TiCl3 VCl3 CrCl3 MnCl3 FeCl3
+4
TiO2 VO2 MnO2
TiCl4 VCl4 CrCl4
+5 V2O5
+6 CrO3
+7 Mn2O7
Oxidation states of the elements of the first transition
series in their oxides and chlorides

67
Oxidation states of the elements of the first transition
series in their compounds
Element Possible oxidation state
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
Element Possible oxidation state
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
+3
+1 +2 +3 +4
+1 +2 +3 +4 +5
+1 +2 +3 +4 +5 +6
+1 +2 +3 +4 +5 +6 +7
+1 +2 +3 +4 +5 +6
+1 +2 +3 +4 +5
+1 +2 +3 +4 +5
+1 +2 +3
+2

68
1. Scandium and zinc do not exhibit
variable oxidation states
• Scandium of the oxidation state +3
 the stable electronic configuration
of argon (i.e. 1s22s22p63s23p6)
• Zinc of the oxidation state +2
 the stable electronic configuration
of [Ar] 3d10

69
2. (a) All elements of the first transition
series (except Sc) can show an
oxidation state of +2
(b) All elements of the first transition
series (except Zn) can show an
oxidation state of +3

70
3. Manganese has the highest oxidation
state +7
E.g. MnO4
-, Mn2O7
Mn7+ ions do not exist.

71
The +7 state of Mn does not mean that
all 3d and 4s electrons are removed
from Mn to give Mn7+.
Instead, Mn forms covalent bonds with
oxygen atoms by making use of its half
filled orbitals
Mn
O
O
O
O-

72
Draw the structure of Mn2O7
Mn
O
O
O
O
Mn
O
O
O

73
3. Manganese has the highest oxidation
state +7
• The highest possible oxidation state
= the total no. of the 3d and 4s electrons
 inner electrons (3s, 3p…) are not
involved in covalent bond formation

74
4. For elements after manganese, there
is a reduction in the number of
possible oxidation states
• The 3d electrons are held more firmly
 the decrease in the number of
unpaired electrons
 the increase in nuclear charge

75
Stability : - Mn2+(aq) > Mn3+(aq)
[Ar] 3d5 [Ar] 3d4
5. The relative stability of various
oxidation states is correlated with the
stability of electronic configurations
o
hydrationH : Fe3+ > Fe2+
Major factor
Major factor
Fe3+(aq) > Fe2+(aq)
[Ar] 3d5 [Ar] 3d6

76
Stability : -
Zn2+(aq) > Zn+(aq)
[Ar] 3d10 [Ar] 3d104s1
5. The relative stability of various
oxidation states is correlated with the
stability of electronic configurations
o
hydrationH : Zn2+ > Zn+
Major factor

77
• The compounds of vanadium, vanadium
 oxidation states of +2, +3, +4 and +5
 forms ions of different oxidation
states
 show distinctive colours in aqueous
solutions
1. Variable Oxidation States of Vanadium and
their Interconversions

78
Ion
Oxidation state of
vanadium in the ion
Colour in
aqueous solution
V2+(aq)
V3+(aq)
VO2+(aq)
VO2
+(aq)
+2
+3
+4
+5
Violet
Green
Blue
Yellow
Colours of aqueous ions of vanadium of
different oxidation states

79
• In an acidic medium
 the vanadium(V) state usually
occurs in the form of VO2
+(aq)
dioxovanadium(V) ion
 the vanadium(IV) state occurs in
the form of VO2+(aq)
oxovanadium(IV) ion

80
• In an alkaline medium
 the stable form of the vanadium(V)
state is
VO3
–(aq), metavanadate(V) or
VO4
3–(aq), orthovanadate(V),
in strongly alkaline medium

81
• Compounds with vanadium in its highest
oxidation state (i.e. +5)
 strong oxidizing agents

82
• Vanadium of its lowest oxidation state
(i.e. +2)
 in the form of V2+(aq)
 strong reducing agent
 easily oxidized when exposed to air

83
• The most convenient starting material
 ammonium metavanadate(V) (NH4VO3)
 a white solid
 the oxidation state of vanadium is +5
• Interconversions of the common
oxidation states of vanadium can be
carried out readily in the laboratory

84
1. Interconversions of Vanadium(V) species
VO2
+(aq) V2O5(s) VO3
(aq) VO4
3(aq)
OH
H+
OH
H+
OH
H+
Yellow orange yellow colourless
Vanadium(V) can exist as cation as well as anion

85
VO2
+(aq) V2O5(s) VO3
(aq) VO4
3(aq)
OH
H+
OH
H+
OH
H+
In acidic medium
In alkaline medium
Amphoteric

86
VO2
+(aq) V2O5(s) VO3
(aq) VO4
3(aq)
OH
H+
OH
H+
OH
H+
In acidic medium
In alkaline medium
Amphoteric
Give the equation for the conversion : V2O5  VO2
+
V2O5(s) + 2H+(aq)  2VO2
+(aq) + H2O(l)

87
VO2
+(aq) V2O5(s) VO3
(aq) VO4
3(aq)
OH
H+
OH
H+
OH
H+
In acidic medium
In alkaline medium
Amphoteric
Give the equation for the conversion : V2O5  VO3

V2O5(s) + 2OH(aq)  2VO3
(aq) + H2O(l)

88
VO2
+(aq) V2O5(s) VO3
(aq) VO4
3(aq)
OH
H+
OH
H+
OH
H+
In acidic medium
In alkaline medium
Give the equation for the conversion : VO3
  VO2
+
VO3
(aq) + 2H+(aq)  VO2
+(aq) + H2O(l)
Amphoteric

89
V5+
H
O
H
H
O
H
H
O
H
H
O
H
VO4
3(aq) + 8H3O+
8H2O
O
H
H
V5+ ions does not exist in water since it undergoes
vigorous hydrolysis to give VO4
3
The reaction is favoured in highly alkaline solution
orthovanadate(V) ion

90
V  VO4
3(aq) orthovanadate(V) ion
Cr  CrO4
2(aq) chromate(VI) ion
Mn  MnO4
(aq) manganate(VII) ion
Draw the structures of VO4
3, CrO4
2 and MnO4

O
Cr
O
O-
O-
O
Mn
O
O
O-

91
V5+
H
O
H
H
O
H
H
O
H
H
O
H
VO3
(aq) + 6H3O+
6H2O
O
H
H
The reaction is favoured in alkaline solution
VO3
 is a polymeric anion like SiO3
2
Metavanadate(V) ion

92
Metavanadate(V) ion, (VO3)n
n

93
V5+
H
O
H
H
O
H
H
O
H
H
O
H
VO2
+(aq) + 4H3O+
4H2O
O
H
H
The reaction is favoured in acidic solution

94
2. The action of zinc powder and
concentrated hydrochloric acid
 vanadium(V) ions can be reduced
sequentially to vanadium(II) ions

95
VO2
+(aq) 
yellow
Zn
conc. HCl
VO2+(aq) 
blue
Zn
conc. HCl
V3+(aq) 
green
Zn
conc. HCl
V2+(aq)
violet

96
(a)
Colours of aqueous solutions of compounds containing
vanadium in four different oxidation states:
(a) +5; (b) +4; (c) +3; (d) +2
(b) (c) (d)
VO2
+(aq) VO2+(aq) V3+(aq) V2+(aq)

97
• The feasibility of the changes in oxidation
state of vanadium
 can be predicted using standard
electrode potentials
Half reaction (V)
Zn2+(aq) + 2e– Zn(s)
VO2
+(aq) + 2H+(aq) + e– VO2+(aq) + H2O(l)
VO2+(aq) + 2H+(aq) + e– V3+(aq) + H2O(l)
V3+(aq) + e– V2+(aq)
–0.76
+1.00
+0.34
–0.26

98
• Under standard conditions
 zinc can reduce
1. VO2
+(aq) to VO2+(aq)
> 0
> 0
> 02. VO2+(aq) to V3+(aq)
3. V3+(aq) to V2+(aq)

99
2 × (VO2
+(aq) + 2H+(aq) + e–
VO2+(aq) + H2O(l)) = +1.00 V
–) Zn2+(aq) + 2e– Zn(s) = –0.76 V
2VO2
+(aq) + Zn(s) + 4H+(aq)
2VO2+(aq) + Zn2+(aq) + 2H2O(l)
= +1.76 V

100
2 × (VO2+(aq) + 2H+(aq) + e–
V3+(aq) + H2O(l)) = +0.34 V
–) Zn2+(aq) + 2e– Zn(s) = –0.76 V
2VO2+(aq) + Zn(s) + 4H+(aq)
2V3+(aq) + Zn2+(aq) + 2H2O(l)
= +1.10 V

101
2 × (V3+(aq) + e– V2+(aq))
= –0.26 V
–) Zn2+(aq) + 2e– Zn(s) = –0.76 V
2V3+(aq) + Zn(s) 2V2+(aq) + Zn2+(aq)
= +0.50 V

102
• Manganese
 show oxidation states of +2, +3, +4,
+5, +6 and +7 in its compounds
2. Variable Oxidation States of Manganese and
• The most common oxidation states
 +2, +4 and +7

103
Ion
Oxidation state of
manganese in the ion
Colour
Mn2+
Mn(OH)3
Mn3+
MnO2
MnO4
3
MnO4
2–
MnO4
–
+2
+3
+3
+4
+5
+6
+7
Very pale pink
Dark brown
Red
Black
Bright blue
Green
Purple
Colours of compounds or ions of manganese in
different oxidation states

104
(a)
differernt oxidation states: (a) +2; (b) +3; (c) +4
(b) (c)
Mn2+(aq) Mn(OH)3(aq) MnO2(s)

105
(e)(d)
differernt oxidation states: (d) +6; (e) +7
MnO4
2–(aq) MnO4
–(aq)

106
• Manganese of the oxidation state +2
 the most stable at pH 0
2. Variable Oxidation States of Manganese and
Mn2+Mn3+
+1.50V
Mn
1.18V
MnO4

+1.51V
MnO2
+1.23V

107
Mn(VII)
Explosive on heating and extremely oxidizing
2KMnO4 K2MnO4 + MnO2 + O2
heat
+7 +6 +42 0
 in ON = 2(+2) = +4
 in ON = (1) + (3) = 4

108
Mn(VII)
 in ON = 6(+2) = +12
 in ON = 4(3) = 12
2 0+4+7
4MnO4
 + 4H+ 4MnO2 + 2H2O + 3O2
light
The reaction is catalyzed by light
Acidified KMnO4(aq) is stored in amber bottle

109
Oxidizing power of Mn(VII) depends on
pH of the solution
In an acidic medium (pH 0)
MnO4
–(aq) + 8H+(aq) + 5e– Mn2+(aq) + 4H2O(l)
= +1.51 V
In a neutral or alkaline medium (up to pH 14)
MnO4
–(aq) + 2H2O(l) + 3e– MnO2(s) + 4OH (aq)
= +0.59 V

110
The reaction does not involve H+(aq) nor OH(aq)
Why is the Eo of MnO4
 MnO4
2 Eo = +0.56V
not affected by pH ?
MnO4
(aq) + e MnO4
2 Eo = +0.56V

111
MnO4
(aq) + e MnO4
2 Eo = +0.56V
When [OH(aq)] > 1M
In an acidic medium (pH 0)
MnO4
–(aq) + 8H+(aq) + 5e– Mn2+(aq) + 4H2O(l)
= +1.51 V
In a neutral or alkaline medium (up to pH 14)
MnO4
–(aq) + 2H2O(l) + 3e– MnO2(s) + 4OH (aq)
= +0.59 V
Under what conditions is the following
conversion favoured?

112
Predict if Mn(VI) Mn(VII) + Mn(IV) is feasible at
(i) pH 0 and (ii) pH 14
At pH 0 (1) 2(3)
3MnO4
2(aq) + 4H+(aq) 2MnO4
(aq) + MnO2(s) + 2H2O(l)
Eo
cell = +1.70V (feasible)
At pH 14 (2) 2(3)
3MnO4
2(aq) + 2H2O(l) 2MnO4
(aq) + MnO2(s) + 4OH(aq)
Eo
cell = +0.04V (much less feasible)
MnO4
2(aq) + 4H+(aq) + 2e MnO2(s) + 2H2O(l) Eo = +2.26V
MnO4
2(aq) + 2H2O(l) + 2e MnO2(s) + 4OH(aq) Eo = +0.60V
MnO4
 + e MnO4
2 Eo = +0.56V
(1)
(2)
(3)
Mn(VI) is unstable in acidic medium

113
Mn(IV) Oxidizing in acidic medium
MnO2(s) + 4H+(aq) + 2e– Mn2+(aq) + 2H2O(l)
= 1.23 V
• Used in the laboratory production of chlorine
MnO2(s) + 4HCl(aq)  MnCl2(aq) + 2H2O(l) + Cl2(g)

114
Mn(IV) Reducing in alkaline medium
• Oxidized to Mn(VI) in alkaline medium
2MnO2 + 4OH + O2  2MnO4
2 + 2H2O

115
MnO2 is oxidized to MnO4
2 in alkaline medium
2MnO2 + 4OH + O2  2MnO4
2 + 2H2O
Suggest a scheme to prepare MnO4
 from MnO2
1. 2MnO2 + 4OH + O2  2MnO4
2 + 2H2O
2. 3MnO4
2 + 4H+  2MnO4
 + MnO2 + 2H2O
3. Filter the resulting mixture to remove MnO2
7B

116
Cu+(aq) + e  Cu(s) Eo = +0.52V
Cu2+(aq) + 2e  Cu(s) Eo = +0.34V
Cu2+(aq) is more stable than Cu+(aq)
The only copper(I) compounds which can be stable
in water are those which are
(i) insoluble (e.g. Cu2O, CuI, CuCl)
(ii) complexed with ligands other than water
e.g. [Cu(NH3)4]+
Cu+(aq) + e  Cu(s)
Under these conditions, [Cu+(aq)] 
 Equil. Position shifts to left

117
Estimation of Cu2+ ions
2Cu2+(aq) + 4I(aq)  2CuI(s) + I2(aq)
I2(aq) + 2S2O3
2(aq)  2I(aq) + S4O6
2(aq)
unknown excess white fixed
standard solution

118
• Another striking feature of the
d-block elements is the formation
of complexes
Formation of Complexes

119
A complex is formed when a central
metal atom or ion is surrounded by
other molecules or ions which form
dative covalent bonds with the central
metal atom or ion.
The molecules or ions that donate lone
pairs of electrons to form the dative
covalent bonds are called ligands.

120
• A ligand
 can be an ion or a molecule having
at least one lone pair of electrons
that can be donated to the central
metal atom or ion to form a dative
covalent bond

121
electrically neutral Ni(CO)4
[Co(H2O)6]3+
positively charged
[Fe(CN)6]3
negatively charged
 Complexes can be

122
 A co-ordination compound is either
a neutral complex e.g. Ni(CO)4
or made of
a complex ion and another ion
e.g. [Co(H2O)6]Cl3  [Co(H2O)6]3+ + 3Cl
K3[Fe(CN)6]  3K+ + [Fe(CN)6]3

123
Criteria for complex formation
2. High charge density of the central
metal ions.
1. Presence of vacant and low-energy 3d,
4s, 4p and 4d orbitals in the metal
atoms or ions to accept lone pairs from
ligands.

124
Diagrammatic representation of the formation of a complex

125
[Co(H2O)6]2+
Co :
 
3d 4s 4p 4d
Co2+ :

3d 4s 4p 4d
sp3d2 hybridisation

The six sp3d2 orbitals accept
six lone pairs from six H2O.
Arranged octahedrally to
minimize repulsion between
dative bonds.

126
1. Complexes with Monodentate Ligands
A ligand that forms one dative covalent
bond only is called a monodentate ligand.
• Examples:
neutral  CO, H2O, NH3
anionic  Cl–, CN–, OH–

128
The transition metal ion is the Lewis acid since it
accepts lone pairs of electrons from the ligands
in forming dative covalent bonds.
The ligand is the Lewis base since it donates a
lone pair of electrons to the transition metal ion
in forming dative covalent bonds.
In the formation of complexes, classify the
transition metal ion and the ligand as a Lewis acid
or base. Explain your answer briefly.

129
What is the oxidation state of the central metal ?
Cr3+ Zn2+

130
Co3+

131
Fe3+ Co2+

132
2. Complexes with Bidentate Ligands
A ligand that can form two dative covalent
bonds with a metal atom or ion is called a
bidentate ligand.
A ligand that can form more than one
dative covalent bond with a central metal
atom or ion is called a chelating ligand.

133
Ethylenediamine
(H2NCH2CH2NH2)
Ethylenediamine (en)
Oxalate (C2O4
2–)
oxalate ion (oxo)
The term chelate is derived from Greek, meaning ‘claw’.
The ligand binds with the metal like the great claw of
the lobster.

134
ethylenediamine
oxalate ion

135
3. Complexes formed by Multidentate Ligands
Ligands that can form more than two
dative covalent bonds to a metal atom
or ion are called multidentate ligands.
Some ligands can form as many as six
bonds to a metal atom or ion.
• Example:
 ethylenediaminetetraacetic acid
(abbreviated as EDTA)

136
ethylenediaminetetraacetate ion
 EDTA forms six dative covalent bonds with
the metal ion through six atoms giving a
very stable complex.
 hexadentate ligand

137
EDTA4
Fe2+
[FeEDTA]2
Structure of the complex ion formed by
iron(II) ions and EDTA
?2

138
Uses of EDTA
1. Determining concentrations of metal ions
by complexometric titrations
e.g. determination of water hardness
2. In chelation therapy for mercury poisoning
and lead poisoning
Poisonous Hg2+ and Pb2+ ions are removed
by forming stable complexes with EDTA.
3. Preparing buffer solutions ( )4aa KtoK 1
4. As preservative to prevent catalytic
oxidation of food by metal ions.

139
The coordination number of the central
metal atom or ion in a complex is the
number of dative covalent bonds formed
by the central metal atom or ion in a
complex.
Complex
The central metal atom
or ion in the complex
Coordination
number
[Ag(NH3)2]+
Ag+ 2
[Cu(NH3)4]2+
Cu2+ 4
[Fe(CN)6]3–
Fe3+ 6

140
4. Nomenclature of Transition Metal
Complexes with Monodentate Ligands
IUPAC conventions
1. (a) For any ionic compound
 the cation is named before the
anion
(b) If the complex is neutral
 the name of the complex is the
name of the compound

141
1. (c) In naming a complex (which may be
neutral, a cation or an anion)
 the ligands are named before
the central metal atom or ion
 the liqands are named in
alphabetical order (prefixes not
counted)
(d) The number of each type of ligands
are specified by the Greek prefixes
1  mono- 2  di 3  tri
4  tetra- 5  penta- 6  hexa-

142
1. (e) The oxidation number of the metal
ion in the complex is indicated
immediately after the name of the
metal using Roman numerals
[CrCl2(H2O)4]Cl
tetraaquadichlorochromium(III) chloride
[CoCl3(NH3)3]
triamminetrichlorocobalt(III)
K3[Fe(CN)6]
potassium hexacyanoferrate(III)

143
2. (a) The root names of anionic ligands
always end in “-o”
CN– cyano
Cl– chloro
Br bromo
I iodo
OH hydroxo
NO2
 nitro
SO4
2 sulphato
H hydrido
(b) The names of neutral ligands are
the names of the molecules
 except NH3, H2O, CO and NO

144
Neutral ligand Name of ligand
Ammonia (NH3)
Water (H2O)
Carbon monoxide (CO)
Nitrogen monoxide (NO)
Ammine
Aqua
Carbonyl
Nitrosyl

145
3. (a) If the complex is anionic
 the suffix “-ate” is added to
the end of the name of the metal,
 followed by the oxidation number
of that metal
tetrachlorocuprate(II) ion[CuCl4]2–
hexacyanoferrate(III) ion[Fe(CN)6]3
tetrachlorocobaltate(II) ion[CoCl4]2
Name of the complexFormula

146
Metal Name in anionic complex
Titanium
Vanadium
Chromium
Manganese
Iron
Cobalt
Nickel
Copper
Zinc
Platinum
Titanate
Vanadate
Chromate
Manganate
Ferrate
Cobaltate
Nickelate
Cuprate
Zincate
Platinate
Names of some common metals in anionic complexes

147
3. (b) If the complex is cationic or neutral
 the name of the metal is unchanged
 followed by the oxidation number
of that metal
triamminetrichlorocobalt(III)[CoCl3(NH3)3]
tetraaquadichlorochromium(III) ion[CrCl2(H2O)4]+
Name of the complexFormula

148
(a) Write the names of the following compounds.
(i) [Fe(H2O)6]Cl2
(ii) [Cu(NH3)4]Cl2
(iii) [PtCl4(NH3)2]
(iv) K2[CoCl4]
(v) [Cr(NH3)4SO4]NO3
(vi) [Co(H2O)2(NH3)3Cl]Cl
(vii) K3[AlF6]

149
Hexaaquairon(II) chloride
Tetraamminecopper(II) chloride
Diamminetetrachloroplatinum(IV)
Potassium tetrachlorocobaltate(II)
Tetraamminesulphatochromium(III) nitrate
(i) [Fe(H2O)6]Cl2
(ii) [Cu(NH3)4]Cl2
(iii) [PtCl4(NH3)2]
(iv) K2[CoCl4]
(v) [Cr(NH3)4SO4]NO3

150
(a) (vi) [Co(H2O)2(NH3)3Cl]Cl
triamminediaquachlorocobalt(II) chloride
(vii) K3[AlF6]
potassium hexafluoroaluminate
Al has a fixed oxidation state (+3)
no need to indicate the oxidation state

151
(b) Write the formulae of the following compounds.
(i) pentaamminechlorocobalt(III) chloride
(ii) Ammonium hexachlorotitanate(IV)
(iii) Tetraaquadihydroxoiron(II)
[Co(NH3)5Cl]Cl2
(NH4)2[TiCl6]
[Fe(H2O)4(OH)2]

152
Coordination number
of the central metal
atom or ion
Shape of complex Example
2
linear
[Ag(NH3)2]+
[Ag(CN)2]–
Stereo-structures of complexes
sp hybridized

153
[Cu(NH3)4]2+
[CuCl4]2–
Square planar
[Zn(NH3)4]2+
[CoCl4]2+
Tetrahedral
4
ExampleShape of complex
Coordination number
atom or ion
sp3
dsp2

154
Tetra-coordinated Complexes
(a) Tetrahedral complexes
 tetrahedral shape
blue
[Co(H2O)6]2+
Octahedral, pink

155
(b) Square planar complexes
 have a square planar structure

156
• Example:

157
Coordination number
atom or ion
Shape of complex Example
6
Octahedral
[Cr(NH3)6]3+
[Fe(CN)6]3–
sp3d2

158
Hexa-coordinated Complexes
• Example:

159
6. Displacement of Ligands and Relative
Stability of Complex Ions
Different ligands have different
tendencies to bind with the metal atom/ion
 ligands compete with one another for
the metal atom/ion.
A stronger ligand can displace a weaker
ligand from a complex.

160
6. Displacement of Ligands and Relative
Stability of Complex Ions
[Fe(H2O)6]2+(aq) + 6CN–(aq)
Hexaaquairon(II) ion
[Fe(CN)6]4–(aq) + 6H2O(l)
Hexacyanoferrate(II) ion
Stronger ligand
Weaker ligand
Reversible reaction
Equilibrium position lies to the right
Kst  1024 mol6 dm18

161
[Ni(H2O)6]2+(aq) + 6NH3(aq)
Hexaaquanickel(II) ion
[Ni(NH3)6]2+(aq) + 6H2O(l)
Hexaamminenickel(II) ion
Stronger ligand
Weaker ligand
The greater the equilibrium constant,
the stronger is the ligand on the LHS and
the more stable is the complex on the RHS
The equilibrium constant is called the
stability constant, Kst

162
Spectrochemical Series
 A partial spectrochemical series
listing of ligands from small Δ to
large Δ is given below.
 I− < Br− < S2− < SCN− < Cl− < NO3
− < N3−
< F− < OH− < C2O4
2− ≈ H2O < NCS− <
CH3CN < py (pyridine) < NH3 < en
(ethylenediamine) < bipy (2,2'-
bipyridine) < phen (1,10-
phenanthroline) < NO2
− < PPh3 < CN−
≈ CO

163
Consider the general equilibrium system below,
[M(H2O)x]m+ + xLn [M(L)x](m-xn)+ + xH2O
xnm
x2
xn)(m
x
st
]][L]O)[[M(H
]][[M(L)
K 

 Units = (mol dm3)-x
Kst measures the stability of the complex, [M(L)x](m-xn)+,
relative to the aqua complex, [M(H2O)x]m+

164
Relative strength of some ligands
bonding with copper(II) ions
monodentate
bidentate
multidentate
TAS Expt 6

165
Equilibrium Kst ((mol dm–3)–n)
[Cu(H2O)4]2+(aq) + 4Cl–(aq)
[CuCl4]2–(aq) + 4H2O(l)
[Cu(H2O)4]2+(aq) + 4NH3(aq)
[Cu(NH3)4]2+(aq) + 4H2O(l)
[Cu(H2O)4]2+(aq) + 2H2NCH2CH2NH2(aq)
[Cu(H2NCH2CH2NH2)2]2+(aq) + 4H2O(l)
[Cu(H2O)4]2+(aq) + EDTA4–(aq)
[CuEDTA]2–(aq) + 4H2O(l)
4.2 × 105
1.1 × 1013
1.0 × 1018.7
1.0 × 1018.8
What is the Kst of the formation of [Cu(H2O)4]2+(aq) ?

166
[Cu(H2O)4]2+ + 4H2O [Cu(H2O)4]2+ + 4H2O
1
]]O)[[Cu(H
]]O)[[Cu(H
K 2
42
2
42
st  


167
Factors affecting the stability of complexes
1. The charge density of the central ion
7.7 × 104
4.5 × 1033
[Co(H2O)6]2+(aq) + 6NH3(aq)
[Co(NH3)6]2+(aq) + 6H2O(l)
[Co(H2O)6]3+(aq) + 6NH3(aq)
[Co(NH3)6]3+(aq) + 6H2O(l)
Kst (mol6 dm18)Equilibrium
≈ 1024
≈ 1031
[Fe(H2O)6]2+(aq) + 6CN–(aq)
[Fe(CN)6]4–(aq) + 6H2O(l)
[Fe(H2O)6]3+(aq) + 6CN–(aq)
[Fe(CN)6]3–(aq) + 6H2O(l)

168
2. The nature of ligands
Ability to form complex : -
CN > NH3 > Cl > H2O
[Zn(CN)4]2 Kst = 5  1016 mol4 dm12
[Zn(NH3)4]2+ Kst = 3.8  109 mol4 dm12
[Cu(NH3)4]2+ Kst = 1.1  1013 mol4 dm12
[CuCl4]2+ Kst = 4.2  105 mol4 dm12

169
3. The pH of the solution
In acidic solution, the ligands are protonated
 lone pairs are not available
 the complex decomposes
[Cu(NH3)4]2+(aq) + 4H2O(l) [Cu(H2O)4]2+(aq) + 4NH3(aq)
NH4
+(aq)
H+(aq)
Equilibrium position shifts to the right

170
Consider the stability constants of the following silver
complexes:
Ag+(aq) + 2Cl–(aq) [AgCl2]–(aq)
Kst = 1.1 × 105 mol–2 dm6
Ag+(aq) + 2NH3(aq) [Ag(NH3)2]+(aq)
Kst = 1.6 × 107 mol–2 dm6
Ag+(aq) + 2CN–(aq) [Ag(CN)2]–(aq)
Kst = 1.0 × 1021 mol–2 dm6
What will be formed when CN–(aq) is added to a
solution of [Ag(NH3)2]+?
[Ag(CN)2](aq) and NH3

171
What will be formed when NH3(aq) is added to a solution
of [Ag(CN)2]–?
No apparent reaction
Consider the stability constants of the following silver
complexes:
Ag+(aq) + 2Cl–(aq) [AgCl2]–(aq)
Kst = 1.1 × 105 mol–2 dm6
Ag+(aq) + 2NH3(aq) [Ag(NH3)2]+(aq)
Kst = 1.6 × 107 mol–2 dm6
Ag+(aq) + 2CN–(aq) [Ag(CN)2]–(aq)
Kst = 1.0 × 1021 mol–2 dm6

172
Fe3+(aq) is too acidic.
FeSO4(aq) is used as the antidote for cyanide poisoning
[Fe(H2O)6]2+(aq) + 6CN(aq) [Fe(CN)6]4 + 6H2O(l)
Kst  1  1024 mol6 dm18
Very stable
[Fe(H2O)6]3+(aq) + H2O(l)
[Fe(H2O)5OH]2+(aq) + H3O+(aq)
Why is Fe2(SO4)3(aq) not used as the antidote ?
Only free CN is poisonous

173
[Cu(H2O)4]2+(aq) + Cl(aq) [Cu(H2O)3Cl]+(aq) + H2O(l)
K1 = 6.3102 mol1 dm3
[Cu(H2O)3Cl]+(aq) + Cl(aq) [Cu(H2O)2Cl2](aq) + H2O(l)
K2 = 40 mol1 dm3
[Cu(H2O)2Cl2](aq) + Cl(aq) [Cu(H2O)Cl3](aq) + H2O(l)
K3 = 5.4 mol1 dm3
[Cu(H2O)Cl3](aq) + Cl(aq) [CuCl4]2(aq) + H2O(l)
K1 = 3.1 mol1 dm3
[Cu(H2O)4]2+(aq) + 4Cl(aq) [CuCl4]2(aq) + 4H2O(l)
Kst = K1  K2  K3  K4 = 4.2  105 mol4 dm12

174
K1 > K2 > K3 > K4
Reasons :
1. Statistical effect
On successive displacement, less water
ligands are available to be displaced.

175
K1 > K2 > K3 > K4
Reasons :
[Cu(H2O)Cl3] Cl repulsion
[Cu(H2O)4]2+ Cl attraction
2. Charge effect
On successive displacement, the Cl
experiences more repulsion from the
complex

176
Formula of copper(II) complex
Colour of the
complex
[Cu(H2O)4]2+
[CuCl4]2–
[Cu(NH3)4]2+
[Cu(H2NCH2CH2NH2)]2+
[Cu(EDTA)]2–
Pale blue
Yellow
Deep blue
Violet
Sky blue
Colours of some copper(II) complexes
The displacement of ligands are usually
accompanied with easily observable colour changes

177
The colours of many gemstones are due to the
presence of small quantities of d-block metal ions
Coloured Ions

178
• Most of the d-block metals
 form coloured compounds
Coloured Ions
 due to the presence of the
incompletely filled d orbitals in the
d-block metal ions
3d10 : Zn2+, Cu+; 3d0 : Sc3+, Ti4+
Which aqueous transition metal ion(s) is/are
not coloured ?

179
Number of unpaired
electrons in 3d
orbitals
d-Block metal
ion
Colour in
aqueous solution
0
Sc3+
Ti4+
Zn2+
Cu+
Colourless
Colourless
Colourless
Colourless
1
Ti3+
V4+
Cu2+
Purple
Blue
Blue
Colours of some d-block metal ions in aqueous solutions

180
Number of unpaired
electrons in 3d
orbitals
d-Block metal
ion
Colour in
aqueous solution
2
V3+
Ni2+
Green
Green
3
V2+
Cr3+
Co2+
Violet
Green
Pink

181
Number of unpaired
electrons in 3d
orbitals
d-Block metal
ion
Colour in
aqueous solution
4
Cr2+
Mn3+
Fe2+
Blue
Violet
Green
5
Mn2+
Fe3+
Very pale pink
Yellow

182
Co2+(aq) Fe3+(aq)Zn2+(aq)

183
Cu2+(aq)Fe2+(aq)Mn2+(aq)

184
A substance absorbs visible light of a certain
wavelength
 reflects or transmits visible light of
other wavelengths (complimentary colour)
 appears coloured
Coloured ion
Light
absorbed
Light reflected or
transmitted
[Cu(H2O)4]2+(aq) Yellow Blue
[CuCl4]2(aq) Blue Yellow

185
Blue
Yellow
Magenta
Green
RedCyan
Violet
Greenish yellow
Complimentary colour
chart
Blue light absorbed
Appears yellow
Yellow light absorbed
Appears blue

186
• The absorption of visible light is due to the
d-d electronic transition
3d  3d
i.e. an electron jumping from a lower 3d
orbital to a higher 3d orbital
Coloured Ions

187
In gaseous state,
the five 3d orbitals are degenerate
i.e. they are of the same energy level
In the presence of ligands,
The five 3d orbitals interact with the
orbitals of ligands and split into two groups
of orbitals with slightly different energy
levels

188
The splitting of the degenerate 3d orbitals of
a d-block metal ion in an octahedral complex
ge
gt2
222
yxz
d,d 
yzxzxy d,d,d
distributes along x and y axes
distributes along z axis
Interact more strongly with
the orbitals of ligands

189
Higher energy eg
22
yx
d 

190
Criterion for d-d transition : -
presence of unpaired d electrons in the d-
block metal atoms or ions
Or presence of incompletely filled d-subshell
d-d transition is possible for 3d1 to 3d9
arrangements
d-d transition is NOT possible for 3d0 & 3d10
arrangements

191
3d9 : d-d transition is possible


Cu2+
H2O as ligand

192


*Cu2+
Yellow light absorbed,
appears blue
H2O as ligand

193

Fe2+

194


*Fe2+
Magenta light absorbed,
appears green

195
3d10 : d-d transition NOT possible

Zn2+


196
3d0 : d-d transition NOT possible
Sc3+

197
E
E depends on
1. the nature and charge of metal ion
[Fe(H2O)6]2+ green,
[Fe(H2O)6]3+ yellow
[Cu(H2O)4]2+ blue,
[CuCl4]2 yellow
2. the nature of ligand

198
Why does Na+(aq) appear colourless ?
Coloured Ions
3d0 : d-d transition is NOT possible
2p  3s transition involves absorption
of radiation in the UV region.

199
• The d-block metals and their compounds
 important catalysts in industry and
biological systems
Catalytic Properties of Transition
Metals and their Compounds

200
d-Block
metal
Catalyst Reaction catalyzed
V
V2O5 or
vanadate(V) (VO3
–)
Contact process
2SO2(g) + O2 (g) 2SO3(g)
Fe Fe
Haber process
N2(g) + 3H2(g) 2NH3(g)
The use of some d-block metals and their compounds as
catalysts in industry

201
d-Block
metal
Catalyst Reaction catalyzed
Ni Ni
Hardening of vegetable oil
(Manufacture of margarine)
RCH = CH2 + H2  RCH2CH3
Pt Pt
Catalytic oxidation of ammonia
(Manufacture of nitric(V) acid)
4NH3(g) + 5O2(g)  4NO(g) + 6H2O(l)
The use of some d-block metals and their compounds as
catalysts in industry

202
• The d-block metals and their compounds
exert their catalytic actions in either
 heterogeneous catalysis
 homogeneous catalysis

203
• Generally speaking, the function of a
catalyst
 provides an alternative reaction
pathway of lower activation energy
 enables the reaction to proceed
faster than the uncatalyzed one

204
1. Heterogeneous Catalysis
• The catalyst and reactants
 exist in different states
• The most common heterogeneous
catalysts
 finely divided solids for gaseous
reactions

205
A heterogeneous catalyst provides a
suitable reaction for the reactants
to come close together and react.

206
• Example:
The synthesis of gaseous ammonia from
nitrogen and hydrogen (i.e. Haber
process)
N2(g) + 3H2(g) 2NH3(g)

207
• In the absence of a catalyst
 the formation of gaseous ammonia
proceeds at an extremely low rate
• The probability of collision of four
gaseous molecules (i.e. one nitrogen and
three hydrogen molecules)
 very small

208
• The four reactant molecules
 collide in proper orientation in order
to form the product
• The bond enthalpy of the reactant (N  N),
 very large
 the reaction has a high activation
energy

209
• In the presence of iron as catalyst
 the reaction proceeds much faster
 provides an alternative reaction
pathway of lower activation energy

210
• Fe is a solid
• H2, N2 and NH3 are gases
• The catalytic action occurs at the interface
between these two states
• The metal provides an active reaction
surface for the reaction to occur

211
1. Gaseous nitrogen and hydrogen
molecules
 diffuse to the surface of the
catalyst
2. The gaseous reactant molecules
 adsorbed (i.e. adhered) on the
surface of the catalyst

212
2. The iron metal
 many 3d electrons and low-lying
vacant 3d orbitals
 form bonds with the reactant
molecules
 adsorb them on its surface
 weakens the bonds present in the
reactant molecules

213
2. The free nitrogen and hydrogen atoms
 come into contact with each other
 readily to react and form the product
3. The weak interaction between the
product and the iron surface
 gaseous ammonia molecules desorb
easily

214
The catalytic mechanism of the formation of
gaseous ammonia from nitrogen and hydrogen

215

216

217

218

219
43.3 Characteristic Properties of the d-Block Elements and their compound
(SB p.162)
• Sometimes, the reactants
 in aqueous or liquid state
• Other example:
The decomposition of hydrogen peroxide
2H2O2(aq)  2H2O(l) + O2(g)
MnO2(s) as the catalyst

220
Energy profiles of the reaction of nitrogen and hydrogen to
form gaseous ammonia in the presence and absence of a
heterogeneous catalyst

221
2. Homogeneous Catalysis
• A homogeneous catalyst
 the same state as the reactants and
products
 the catalyst forms an intermediate
with the reactants in the reaction
 changes the reaction mechanism to
an another one with a lower
activation energy

222
In homogeneous catalysis, the ability of
the d-block metals to exhibit variable
oxidation states enables the formation of
the reaction intermediates.
• Example:
The reaction between peroxodisulphate(VI)
ions (S2O8
2–) and iodide ions (I–)

223
• Peroxodisulphate(VI) ions
 oxidize iodide ions to iodine in an
aqueous solution
 themselves being reduced to
sulphate(VI) ions
S2O8
2–(aq) + 2I–(aq)
2SO4
2–(aq) + I2 (aq)
V.Eo
cell 511

224
• Iron(III) ions
 take part in the reaction by oxidizing
iodide ions to iodine
 themselves being reduced to iron(II)
ions
2I–(aq) + 2Fe3+(aq)
I2(aq) + 2Fe2+(aq) = +0.23 V
• The reaction is very slow due to strong
repulsion between like charges.

225
• Iron(II) ions
 subsequently oxidized by
peroxodisulphate(VI) ion
 the original iron(III) ions are
regenerated
2Fe2+(aq) + S2O8
2–(aq)
2Fe3+(aq) + 2SO4
2–(aq) = +1.28 V

226
• The overall reaction:
2I–(aq) + 2Fe3+(aq)
I2(aq) + 2Fe2+(aq) = +0.23 V
S2O8
2–(aq) + 2I–(aq)
2SO4
2–(aq) + I2(aq) = +1.51 V
2Fe2+(aq) + S2O8
2–(aq)
+) 2Fe3+(aq) + 2SO4
2–(aq) = +1.28 V
Feasible reaction

227
43.3 Characteristic Properties of the d-Block Elements and their compound
(SB p.164)
• Iron(III) ions
 catalyze the reaction
 acting as an intermediate for the
transfer of electrons between
peroxodisulphate(VI) ions and iodide
ions

228
• Iodide ions
 reduce Fe3+ to Fe2+
• Peroxodisulphate(VI) ions
 oxidize Fe2+ to Fe3+

The Transition Metals: Properties and Applications

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