Transcript d_block - ilc.edu.hk

The d-Block Elements
1
Introduction
•
d-block elements
 locate between the s-block and
p-block
 known as transition elements
 occur in the fourth and subsequent
periods of the Periodic Table
2
d-block elements
period 4
period 5
period 6
period 7
3
Introduction
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
4
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 imcomplete.
Sc  Sc3+ 3d0
5
Zn  Zn2+ 3d10
Introduction
Cu
Cu+ 3d10
not transitional
Cu2+ 3d9
transitional
6
The first transition series
the first horizontal row of the d-block elements
7
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)
8
Characteristics of transition elements
4. Formation of coloured compounds/ions
(cf. colourless ions of s-block elements)
5. Formation of complexes
6. Catalytic properties
9
Electronic Configurations
The building up of electronic configurations
of elements follow:
 Aufbau principle
 Pauli exclusion principle
 Hund’s rule
10
Electronic Configurations
11
•
3d and 4s sub-shells are very close to
each other in energy.
•
Relative energy of electrons in subshells depends on the effective nuclear
charge they experience.
•
Electrons enter 4s sub-shell first
•
Electrons leave 4s sub-shell first
Cu
Cu2+
Relative energy levels of orbitals
in atom and in ion
12
Electronic Configurations
•
Valence electrons in the inner 3d orbitals
•
Examples:
 The electronic configuration of
scandium: 1s22s22p63s23p63d14s2
 The electronic configuration of zinc:
1s22s22p63s23p63d104s2
13
Electronic configurations of the first series of the
d-block elements
14
Element
Atomic number
Electronic configuration
Scandium
21
[Ar] 3d 14s2
Titanium
22
[Ar] 3d 24s2
Vanadium
23
[Ar] 3d 34s2
Chromium
24
[Ar] 3d 54s1
Manganese
25
[Ar] 3d 54s2
Iron
26
[Ar] 3d 64s2
Cobalt
27
[Ar] 3d 74s2
Nickel
28
[Ar] 3d 84s2
Copper
29
[Ar] 3d 104s1
Zinc
30
[Ar] 3d 104s2
•
A half-filled or fully-filled d sub-shell
has extra stability
15
d -Block Elements as Metals
•
d-Block elements are typical metals
Physical properties of d-Block elements :
 good conductors of heat and electricity
 hard and strong
 malleable and ductile
16
d -Block Elements as Metals
•
Physical properties of d-Block elements:
 lustrous
 high melting points and boiling points
•
Exceptions : Mercury
 low melting point
 liquid at room temperature and
pressure
17
d -Block Elements as Metals
•
d-block elements
 extremely useful as construction
materials
 strong and unreactive
18
d -Block Elements as Metals
•
Iron
 used for construction and making
machinery nowadays
 abundant
 easy to extract
19
cheap
d -Block Elements as Metals
•
Iron
 corrodes easily
 often combined with other
elements to form steel
 harder and more resistant to
corrosion
20
d -Block Elements as Metals
•
Titanium
Corrosion resistant, light, strong and
withstand large temperature changes
 used to make aircraft and space
shuttles
 expensive
21
d -Block Elements as Metals
•
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
22
d -Block Elements as Metals
•
Chromium
 confers inertness to stainless steel
•
Manganese
confers hardness & wearing resistance to
its alloys
e.g. duralumin : alloy of Al with Mn/Mg/Cu
23
Atomic Radii and Ionic Radii
•
Two features can be observed:
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
1.
24
Atomic Radii and Ionic Radii
Variation in atomic radius
of the first 36 elements
25
26
27
On moving across the Period,
(i) Nuclear charge 
(ii) Shielding effect (repulsion between e-) 
(i)  (ii)
(i) > (ii)
28
(ii) > (i)
Atomic Radii and Ionic Radii
•
At the beginning of the series
 atomic number 
 effective nuclear charge 
 the electron clouds are pulled
closer to the nucleus
 atomic size 
29
•
In the middle of the series
 more electrons enter the inner
3d sub-shell
 The inner 3d electrons shield the
outer 4s electrons effectively
 the effective nuclear charge
experienced by 4s electrons increases
very slowly
 only a slow decrease in atomic radius
in this region
30
Atomic Radii and Ionic Radii
•
At the end of the series
 the screening and repulsive effects
of the electrons in the 3d subshell become even stronger
 Atomic size 
31
Comparison of Some Physical and
Chemical Properties between the
d-Block and s-Block Elements
•
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
32
1. Density
Densities (in g cm–3) of the s-block elements and
the first series of the d-block elements at 20C
33
1. Density
•
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
34
1. Density
•
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
35
2. Ionization Enthalpy
K  Ca (sharp ) ; Ca  Sc (slight )
Ionization enthalpy (kJ mol–1)
Element
36
1st
2nd
3rd
4th
K
418
3 070
4 600
5 860
Ca
590
1 150
4 940
6 480
Sc
632
1 240
2 390
7 110
Ti
661
1 310
2 720
4 170
V
648
1 370
2 870
4 600
Cr
653
1 590
2 990
4 770
2. Ionization Enthalpy
Sc  Cu (slight ) ; Cu  Zn (sharp )
Ionization enthalpy (kJ mol–1)
Element
37
1st
2nd
3rd
4th
Cr
653
1 590
2 990
4 770
Mn
716
1 510
3 250
5 190
Fe
762
1 560
2 960
5 400
Co
757
1 640
3 230
5 100
Ni
736
1 750
3 390
5 400
Cu
745
1 960
3 550
5 690
Zn
908
1 730
3 828
5 980
2. Ionization Enthalpy
•
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
38
Sharp  across periods 1, 2 and 3
Slight  across the transition series
39
2. Ionization Enthalpy
•
Going across the first transition series
 the nuclear charge of the elements
increases
 additional electrons are added to
the ‘inner’ 3d sub-shell
40
2. Ionization Enthalpy
•
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.
41
Electron has to be removed from
completely filled 3p subshell
3d5
3d5
3d5
42
Cr+
Fe3+
Mn2+
3d10
d10/s2
2. Ionization Enthalpy
•
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
43
3. Melting Points and Hardness
d-block >> s-block

1. both 4s and 3d e- are involved in the
formation of metal bonds
2. d-block atoms are smaller
1541 1668 1910 1907 1246 1538 1495 1455 1084 419
44
3. Melting Points and Hardness
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
45
Sc
Ti
V
Cr
1541 1668 1910 1907
Mn
Fe
Co
Ni
Cu
Zn
1246 1538 1495 1455 1084 419
 Unpaired electrons are relatively
more involved in the sea of electrons
46
Sc
Ti
V
Cr
1541 1668 1910 1907
Sc

Ti


V


Mn
Fe
Co
Ni
Cu
Zn
1246 1538 1495 1455 1084 419
3d
4s




1. m.p.  from Sc to V due to the  of
unpaired d-electrons (from d1 to d3)
47
Sc
Ti
V
Cr
1541 1668 1910 1907
Mn
Co
Ni
Cu
Zn
1246 1538 1495 1455 1084 419
3d
Fe

Co
 
Ni
  

Fe
4s











2. m.p.  from Fe to Zn due to the 
of unpaired d-electrons (from 4 to 0)
48
Sc
Ti
V
Cr
1541 1668 1910 1907
Mn
Fe
Co
Ni
Cu
Zn
1246 1538 1495 1455 1084 419
3. Cr has the highest no. of unpaired
electrons but its m.p. is lower than V.
3d
Cr



4s



It is because the electrons in the
half-filled d-subshell are relatively
less involved in the sea of electrons.
49
Sc
Ti
V
Cr
1541 1668 1910 1907
Mn
Fe
Co
Ni
Cu
Zn
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.
50
3. Melting Points and Hardness
•
The hardness of a metal depends on
 the strength of the metallic bonds
•
The metallic bonds of the d-block
elements are stronger than those of the
s-block elements
 much harder than the s-block
elements
51
Mohs scale : - A measure of hardness
Talc
0
K
Diamond
10
Ca
Sc
Ti
V
Cr
0.5 1.5
3.0
4.5
6.1
9.0 5.0
52
Mn
Fe
Co
Ni
Cu
Zn
4.5
--
--
2.8
2.5
4. Reaction with Water
•
In general, the s-block elements
 react vigorously with water to form
metal hydroxides and hydrogen
•
The d-block elements
 react very slowly with cold water
 react with steam to give metal oxides
and hydrogen
53
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)
54
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+
55
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
56
Variable Oxidation States
•
One of the most striking properties
 variable oxidation states
•
The 3d and 4s electrons are
 in similar energy levels
 available for bonding
57
Variable Oxidation States
•
Elements of the first transition series
 form ions of roughly the same
stability by losing different
numbers of the 3d and 4s electrons
58
Oxidation states of the elements of the first transition
series in their oxides and chlorides
Oxidation
Oxides / Chloride
states
Cu2O
+1
+2
+3
+4
+5
+6
+7
59
Cu2Cl2
TiO
VO
CrO
MnO
FeO
TiCl2
VCl2
CrCl2 MnCl2 FeCl2 CoCl2 NiCl2
Sc2O3 Ti2O3 V2O3
Cr2O3 Mn2O3 Fe2O3
ScCl3 TiCl3
VCl3
CrCl3 MnCl3 FeCl3
TiO2
VO2
MnO2
TiCl4
VCl4
CrCl4
V2O5
CrO3
Mn2O7
CoO
NiO
Ni2O3 • xH2O
CuO
ZnO
CuCl2 ZnCl2
Oxidation states of the elements of the first transition
series in their compounds
Element
Possible oxidation state
Sc
Ti
+1
+2
+3
+4
V
+1
+2
+3
+4
+5
Cr
+1
+2
+3
+4
+5
+6
Mn
+1
+2
+3
+4
+5
+6
Fe
+1
+2
+3
+4
+5
+6
Co
+1
+2
+3
+4
+5
Ni
+1
+2
+3
+4
+5
Cu
+1
+2
+3
Zn
60
+3
+2
+7
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
61
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
62
3. Manganese has the highest oxidation
state +7
E.g. MnO4-, Mn2O7
Mn7+ ions do not exist.
63
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
O
filled orbitals
Mn
O
O
O
64
-
Draw the structure of Mn2O7
O
O
Mn
Mn
O
O
O
65
O
O
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
66
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
67
5. The relative stability of various
oxidation states is correlated with the
stability of electronic configurations
Stability : - Mn2+(aq)
Major factor
>
[Ar] 3d5
Fe3+(aq)
[Ar] 3d4
>
[Ar] 3d5
Major factor
68
 H
Mn3+(aq)
o
:
hydration
Fe2+(aq)
[Ar] 3d6
Fe3+ > Fe2+
5. The relative stability of various
oxidation states is correlated with the
stability of electronic configurations
Stability : -
Zn2+(aq)
>
[Ar] 3d10
Major factor
69
 H
o
:
hydration
Zn+(aq)
[Ar] 3d104s1
Zn2+ > Zn+
1. Variable Oxidation States of Vanadium and
their Interconversions
•
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
70
Colours of aqueous ions of vanadium of
different oxidation states
Ion
Oxidation state of
Colour in
vanadium in the ion aqueous solution
V2+(aq)
+2
Violet
V3+(aq)
+3
Green
VO2+(aq)
+4
Blue
VO2+(aq)
+5
Yellow
71
1. Variable Oxidation States of Vanadium and
their Interconversions
•
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
72
1. Variable Oxidation States of Vanadium and
their Interconversions
•
In an alkaline medium
 the stable form of the vanadium(V)
state is
VO3–(aq), metavanadate(V) or
VO43–(aq), orthovanadate(V),
in strongly alkaline medium
73
1. Variable Oxidation States of Vanadium and
their Interconversions
•
Compounds with vanadium in its highest
oxidation state (i.e. +5)
 strong oxidizing agents
74
1. Variable Oxidation States of Vanadium and
their Interconversions
•
Vanadium of its lowest oxidation state
(i.e. +2)
 in the form of V2+(aq)
 strong reducing agent
 easily oxidized when exposed to air
75
1. Variable Oxidation States of Vanadium and
their Interconversions
•
Interconversions of the common
oxidation states of vanadium can be
carried out readily in the laboratory
•
The most convenient starting material
 ammonium metavanadate(V) (NH4VO3)
 a white solid
 the oxidation state of vanadium is +5
76
1. Variable Oxidation States of Vanadium and
their Interconversions
1. Interconversions of Vanadium(V) species
VO2+(aq)
Yellow
OH
H+
V2O5(s)
orange
OH
H+
VO3(aq)
yellow
OH
H+
VO43(aq)
colourless
Vanadium(V) can exist as cation as well as anion
77
1. Variable Oxidation States of Vanadium and
their Interconversions
1. Interconversions of Vanadium(V) species
VO2+(aq)
Yellow
OH
H+
In acidic medium
78
V2O5(s)
orange
OH
H+
Amphoteric
VO3(aq)
yellow
OH
H+
VO43(aq)
colourless
In alkaline medium
1. Variable Oxidation States of Vanadium and
their Interconversions
1. Interconversions of Vanadium(V) species
VO2+(aq)
Yellow
OH
H+
In acidic medium
V2O5(s)
orange
OH
H+
Amphoteric
VO3(aq)
yellow
OH
H+
VO43(aq)
colourless
In alkaline medium
Give the equation for the conversion : V2O5  VO2+
79
V2O5(s) + 2H+(aq)  2VO2+(aq) + H2O(l)
1. Variable Oxidation States of Vanadium and
their Interconversions
1. Interconversions of Vanadium(V) species
VO2+(aq)
Yellow
OH
H+
In acidic medium
V2O5(s)
orange
OH
H+
Amphoteric
VO3(aq)
yellow
OH
H+
VO43(aq)
colourless
In alkaline medium
Give the equation for the conversion : V2O5  VO3
80
V2O5(s) + 2OH(aq)  2VO3(aq) + H2O(l)
1. Variable Oxidation States of Vanadium and
their Interconversions
1. Interconversions of Vanadium(V) species
VO2+(aq)
Yellow
OH
H+
In acidic medium
V2O5(s)
orange
OH
H+
Amphoteric
VO3(aq)
yellow
OH
H+
VO43(aq)
colourless
In alkaline medium
Give the equation for the conversion : VO3  VO2+
81
VO3(aq) + 2H+(aq)  VO2+(aq) + H2O(l)
H
H
H
O
H
H
V5+
O
O
H
H
H
O
H
O
H
8H2O
orthovanadate(V) ion VO43(aq)
+ 8H3O+
V5+ ions does not exist in water since it undergoes
vigorous hydrolysis to give VO43
The reaction is favoured in highly alkaline solution
82
V  VO43(aq) orthovanadate(V) ion
Cr  CrO42(aq) chromate(VI) ion
Mn  MnO4(aq) manganate(VII) ion
Draw the structures of VO43, CrO42 and MnO4
O
O
Cr
O
83
O
O-
-
Mn
O
OO
H
H
H
O
H
H
V5+
O
O
H
H
H
O
H
O
H
6H2O
Metavanadate(V) ion VO3(aq)
+ 6H3O+
The reaction is favoured in alkaline solution
VO3 is a polymeric anion like SiO32
84
Metavanadate(V) ion, (VO3)nn
85
H
H
H
O
H
H
V5+
O
O
H
H
H
O
H
O
H
4H2O
VO2+(aq) + 4H3O+
The reaction is favoured in acidic solution
86
1. Variable Oxidation States of Vanadium and
their Interconversions
2. The action of zinc powder and
concentrated hydrochloric acid
 vanadium(V) ions can be reduced
sequentially to vanadium(II) ions
87
1. Variable Oxidation States of Vanadium and
their Interconversions
VO2
+(aq)
yellow

Zn
conc. HCl
VO2+(aq)

Zn
conc. HCl
blue
V3+(aq)
green
88
 V2+(aq)
Zn
conc. HCl
violet
(a)
(b)
VO2+(aq)
VO2+(aq)
(c)
(d)
V3+(aq)
V2+(aq)
Colours of aqueous solutions of compounds containing
vanadium in four different oxidation states:
(a) +5; (b) +4; (c) +3; (d) +2
89
•
The feasibility of the changes in oxidation
state of vanadium
 can be predicted using standard
electrode potentials
Half reaction
Zn2+(aq) + 2e–
(V)
–0.76
Zn(s)
VO2+(aq) + 2H+(aq) + e–
VO2+(aq) + H2O(l)
+1.00
VO2+(aq) + 2H+(aq) + e–
V3+(aq) + H2O(l)
+0.34
V3+(aq) + e–
90
V2+(aq)
–0.26
1. Variable Oxidation States of Vanadium and
their Interconversions
•
Under standard conditions
 zinc can reduce
91
1. VO2+(aq) to VO2+(aq)
>0
2. VO2+(aq) to V3+(aq)
>0
3. V3+(aq) to V2+(aq)
>0
1. Variable Oxidation States of Vanadium and
their Interconversions
2 × (VO2+(aq) + 2H+(aq) + e–
VO2+(aq) + H2O(l))
–) Zn2+(aq) + 2e–
Zn(s)
= +1.00 V
= –0.76 V
2VO2+(aq) + Zn(s) + 4H+(aq)
2VO2+(aq) + Zn2+(aq) + 2H2O(l)
= +1.76 V
92
1. Variable Oxidation States of Vanadium and
their Interconversions
2 × (VO2+(aq) + 2H+(aq) + e–
V3+(aq) + H2O(l))
–) Zn2+(aq) + 2e–
Zn(s)
= +0.34 V
= –0.76 V
2VO2+(aq) + Zn(s) + 4H+(aq)
2V3+(aq) + Zn2+(aq) + 2H2O(l)
= +1.10 V
93
1. Variable Oxidation States of Vanadium and
their Interconversions
2 × (V3+(aq) + e–
–) Zn2+(aq) + 2e–
2V3+(aq) + Zn(s)
V2+(aq))
= –0.26 V
Zn(s)
= –0.76 V
2V2+(aq) + Zn2+(aq)
= +0.50 V
94
2. Variable Oxidation States of Manganese and
their Interconversions
•
Manganese
 show oxidation states of +2, +3, +4,
+5, +6 and +7 in its compounds
•
The most common oxidation states
 +2, +4 and +7
95
Colours of compounds or ions of manganese in
different oxidation states
96
Ion
Oxidation state of
manganese in the ion
Colour
Mn2+
+2
Very pale pink
Mn(OH)3
+3
Dark brown
Mn3+
+3
Red
MnO2
+4
Black
MnO43
+5
Bright blue
MnO42–
+6
Green
MnO4–
+7
Purple
(a)
(b)
Mn2+(aq)
Mn(OH)3(aq)
(c)
MnO2(s)
Colours of compounds or ions of manganese in
differernt oxidation states: (a) +2; (b) +3; (c) +4
97
(d)
MnO42–(aq)
(e)
MnO4–(aq)
Colours of compounds or ions of manganese in
differernt oxidation states: (d) +6; (e) +7
98
2. Variable Oxidation States of Manganese and
their Interconversions
•
Manganese of the oxidation state +2
 the most stable at pH 0
Mn3+
+1.50V
+1.51V
MnO4
99
Mn2+
1.18V
+1.23V
MnO2
Mn
Mn(VII)
Explosive on heating and extremely oxidizing
+7 2
2KMnO4
heat
+4
0
K2MnO4 + MnO2 + O2
 in ON = 2(+2) = +4
 in ON = (1) + (3) = 4
100
+6
Mn(VII)
+7 2
4MnO4 + 4H+
light
+4
0
4MnO2 + 2H2O + 3O2
 in ON = 6(+2) = +12
 in ON = 4(3) = 12
The reaction is catalyzed by light
Acidified KMnO4(aq) is stored in amber bottle
101
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
102
Why is the Eo of MnO4
MnO42 Eo = +0.56V
not affected by pH ?
MnO4(aq) + e
MnO42
Eo = +0.56V
The reaction does not involve H+(aq) nor OH(aq)
103
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?
MnO4(aq) + e
MnO42 Eo = +0.56V
When [OH(aq)] > 1M
104
Predict if Mn(VI)
Mn(VII) + Mn(IV) is feasible at
(i) pH 0 and (ii) pH 14
(1) MnO42(aq) + 4H+(aq) + 2e
MnO2(s) + 2H2O(l) Eo = +2.26V
(2) MnO42(aq) + 2H2O(l) + 2e
MnO2(s) + 4OH(aq) Eo = +0.60V
(3) MnO4 + e
MnO42
Eo = +0.56V
At pH 0 (1) 2(3)
3MnO42(aq) + 4H+(aq)
Eocell = +1.70V (feasible)
At pH 14 (2) 2(3)
3MnO42(aq) + 2H2O(l)
2MnO4(aq) + MnO2(s) + 2H2O(l)
Mn(VI) is unstable in acidic medium
2MnO4(aq) + MnO2(s) + 4OH(aq)
Eocell = +0.04V (much less feasible)
105
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)
106
Mn(IV) Reducing in alkaline medium
•
Oxidized to Mn(VI) in alkaline medium
2MnO2 + 4OH + O2  2MnO42 + 2H2O
107
MnO2 is oxidized to MnO42 in alkaline medium
2MnO2 + 4OH + O2  2MnO42 + 2H2O
Suggest a scheme to prepare MnO4 from MnO2
1. 2MnO2 + 4OH + O2  2MnO42 + 2H2O
2. 3MnO42 + 4H+  2MnO4 + MnO2 + 2H2O
3. Filter the resulting mixture to remove MnO2
7B
108
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)] 
109
 Equil. Position shifts to left
Estimation of Cu2+ ions
2Cu2+(aq) + 4I(aq)  2CuI(s) + I2(aq)
unknown
excess
white
fixed
I2(aq) + 2S2O32(aq)  2I(aq) + S4O62(aq)
standard solution
110
Formation of Complexes
•
111
Another striking feature of the
d-block elements is the formation
of complexes
Formation of Complexes
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.
112
Formation of Complexes
•
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
113
Formation of Complexes
 Complexes can be
electrically neutral Ni(CO)4
positively charged
[Co(H2O)6]3+
negatively charged [Fe(CN)6]3
114
 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
115
Criteria for complex formation
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.
2. High charge density of the central
metal ions.
116
Diagrammatic representation of the formation of a complex
117
[Co(H2O)6]2+
Co :
3d
  
Co2+ :
4s


3d
  
4d
4p
4d

4s

4p

The six sp3d2 orbitals accept
six lone pairs from six H2O.
Arranged octahedrally to
minimize repulsion between
dative
bonds.
118
sp3d2 hybridisation
     
1. Complexes with Monodentate Ligands
A ligand that forms one dative covalent
bond only is called a monodentate ligand.
•
119
Examples:
neutral 
CO, H2O, NH3
anionic 
Cl–, CN–, OH–
120
In the formation of complexes, classify the
transition metal ion and the ligand as a Lewis acid
or base. Explain your answer briefly.
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.
121
What is the oxidation state of the central metal ?
Cr3+
122
Zn2+
What is the oxidation state of the central metal ?
Co3+
123
What is the oxidation state of the central metal ?
Fe3+
124
Co2+
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.
125
Ethylenediamine
(H2NCH2CH2NH2)
Oxalate (C2O42–)
ethylenediamine
oxalate ion
The term chelate is derived from Greek, meaning ‘claw’.
The ligand binds with the metal like the great claw of
the lobster.
126
ethylenediamine
127
oxalate ion
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)
128
 EDTA forms six dative covalent bonds with
the metal ion through six atoms giving a
very stable complex.
 hexadentate ligand
ethylenediaminetetraacetate ion
129
Fe2+
EDTA4
2
?
[FeEDTA]2
Structure of the complex ion formed by
iron(II) ions and EDTA
130
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 ( K a to K a )
4.
As preservative to prevent catalytic
oxidation of food by metal ions.
131
1
4
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 Coordination
or ion in the complex
number
[Ag(NH3)2]+
Ag+
2
[Cu(NH3)4]2+
Cu2+
4
[Fe(CN)6]3–
Fe3+
6
132
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
133
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
134
1  mono-
2  di
3  tri
4  tetra-
5  penta-
6  hexa-
1. (e) The oxidation number of the metal
ion in the complex is indicated
immediately after the name of the
metal using Roman numerals
K3[Fe(CN)6]
potassium hexacyanoferrate(III)
[CrCl2(H2O)4]Cl
tetraaquadichlorochromium(III) chloride
[CoCl3(NH3)3]
triamminetrichlorocobalt(III)
135
2. (a) The root names of anionic ligands
always end in “-o”
CN–
cyano
OH
hydroxo
Cl–
chloro
NO2
nitro
Br
bromo
SO42
sulphato
I
iodo
H
hydrido
(b) The names of neutral ligands are
the names of the molecules
 except NH3, H2O, CO and NO
136
137
Neutral ligand
Ammonia (NH3)
Name of ligand
Water (H2O)
Aqua
Carbon monoxide (CO)
Carbonyl
Nitrogen monoxide (NO)
Nitrosyl
Ammine
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
138
Formula
Name of the complex
[CoCl4]2
tetrachlorocobaltate(II) ion
[Fe(CN)6]3
hexacyanoferrate(III) ion
[CuCl4]2–
tetrachlorocuprate(II) ion
Names of some common metals in anionic complexes
139
Metal
Name in anionic complex
Titanium
Titanate
Vanadium
Vanadate
Chromium
Chromate
Manganese
Manganate
Iron
Ferrate
Cobalt
Cobaltate
Nickel
Nickelate
Copper
Cuprate
Zinc
Zincate
Platinum
Platinate
3. (b) If the complex is cationic or neutral
 the name of the metal is unchanged
 followed by the oxidation number
of that metal
140
Formula
Name of the complex
[CrCl2(H2O)4]+
tetraaquadichlorochromium(III) ion
[CoCl3(NH3)3]
triamminetrichlorocobalt(III)
(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]
141
(i) [Fe(H2O)6]Cl2
Hexaaquairon(II) chloride
(ii) [Cu(NH3)4]Cl2
Tetraamminecopper(II) chloride
(iii) [PtCl4(NH3)2]
Diamminetetrachloroplatinum(IV)
(iv) K2[CoCl4]
Potassium tetrachlorocobaltate(II)
(v) [Cr(NH3)4SO4]NO3
Tetraamminesulphatochromium(III) nitrate
142
(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
143
(b) Write the formulae of the following compounds.
(i) pentaamminechlorocobalt(III) chloride
[Co(NH3)5Cl]Cl2
(ii) Ammonium hexachlorotitanate(IV)
(NH4)2[TiCl6]
(iii) Tetraaquadihydroxoiron(II)
[Fe(H2O)4(OH)2]
144
Stereo-structures of complexes
Coordination number
of the central metal
atom or ion
Shape of complex
Example
sp hybridized
[Ag(NH3)2]+
2
[Ag(CN)2]–
linear
145
Stereo-structures of complexes
Coordination number
of the central metal
atom or ion
Shape of complex
sp3
Example
[Zn(NH3)4]2+
[CoCl4]2+
4
Tetrahedral
dsp2
[Cu(NH3)4]2+
[CuCl4]2–
Square planar
146
Tetra-coordinated Complexes
(a) Tetrahedral complexes
 tetrahedral shape
147
[Co(H2O)6]2+
Octahedral, pink
blue
Tetra-coordinated Complexes
(b) Square planar complexes
 have a square planar structure
148
Tetra-coordinated Complexes
•
149
Example:
Stereo-structures of complexes
Coordination number
of the central metal
atom or ion
Shape of complex
Example
sp3d2
[Cr(NH3)6]3+
6
[Fe(CN)6]3–
Octahedral
150
Hexa-coordinated Complexes
•
151
Example:
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.
152
6. Displacement of Ligands and Relative
Stability of Complex Ions
Stronger ligand
[Fe(H2O)6]2+(aq) + 6CN–(aq)
Hexaaquairon(II) ion
Weaker ligand
[Fe(CN)6]4–(aq) + 6H2O(l)
Hexacyanoferrate(II) ion
Reversible reaction
Kst  1024 mol6 dm18
Equilibrium position lies to the right
153
Stronger ligand
[Ni(H2O)6]2+(aq) + 6NH3(aq)
Hexaaquanickel(II) ion
Weaker ligand
[Ni(NH3)6]2+(aq) + 6H2O(l)
Hexaamminenickel(II) ion
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
154
Consider the general equilibrium system below,
[M(H2O)x]m+ + xLn
[M(L)x](m-xn)+ + xH2O
[[M(L)x ]
]
Kst 
m
n x
[[M(H2O)x ] ][L ]
(m xn) 
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+
155
monodentate
bidentate
multidentate
TAS Expt 6
156
Relative strength of some ligands
bonding with copper(II) ions
Equilibrium
[Cu(H2O)4]2+(aq) + 4Cl–(aq)
Kst ((mol dm–3)–n)
4.2 × 105
[CuCl4]2–(aq) + 4H2O(l)
[Cu(H2O)4]2+(aq) + 4NH3(aq)
1.1 × 1013
[Cu(NH3)4]2+(aq) + 4H2O(l)
[Cu(H2O)4]2+(aq) + 2H2NCH2CH2NH2(aq)
1.0 × 1018.7
[Cu(H2NCH2CH2NH2)2]2+(aq) + 4H2O(l)
[Cu(H2O)4]2+(aq) + EDTA4–(aq)
1.0 × 1018.8
[CuEDTA]2–(aq) + 4H2O(l)
What is the Kst of the formation of [Cu(H2O)4]2+(aq) ?
157
[Cu(H2O)4]2+ + 4H2O
[Cu(H2O)4]2+ + 4H2O
[[Cu(H2O)4 ]2 ]
Kst 
1
2
[[Cu(H2O)4 ] ]
158
Factors affecting the stability of complexes
1. The charge density of the central ion
Equilibrium
[Co(H2O)6]2+(aq) + 6NH3(aq)
Kst (mol6 dm18)
7.7 × 104
[Co(NH3)6]2+(aq) + 6H2O(l)
[Co(H2O)6]3+(aq) + 6NH3(aq)
4.5 × 1033
[Co(NH3)6]3+(aq) + 6H2O(l)
[Fe(H2O)6]2+(aq) + 6CN–(aq)
≈ 1024
[Fe(CN)6]4–(aq) + 6H2O(l)
[Fe(H2O)6]3+(aq) + 6CN–(aq)
≈ 1031
[Fe(CN)6]3–(aq) + 6H2O(l)
159
Factors affecting the stability of complexes
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
160
Factors affecting the stability of complexes
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)
H+(aq)
Equilibrium position shifts to the right
161
NH4+(aq)
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
162
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 NH3(aq) is added to a solution
of [Ag(CN)2]–?
No apparent reaction
163
FeSO4(aq) is used as the antidote for cyanide poisoning
[Fe(H2O)6]2+(aq) + 6CN(aq)
Kst  1  1024 mol6 dm18
[Fe(CN)6]4 + 6H2O(l)
Very stable
Only free CN is poisonous
Why is Fe2(SO4)3(aq) not used as the antidote ?
Fe3+(aq) is too acidic.
[Fe(H2O)6]3+(aq) + H2O(l)
[Fe(H2O)5OH]2+(aq) + H3O+(aq)
164
[Cu(H2O)4]2+(aq) + Cl(aq)
[Cu(H2O)3Cl]+(aq) + H2O(l)
[Cu(H2O)3Cl]+(aq) + Cl(aq)
[Cu(H2O)2Cl2](aq) + H2O(l)
[Cu(H2O)2Cl2](aq) + Cl(aq)
K3 = 5.4 mol1 dm3
[Cu(H2O)Cl3](aq) + H2O(l)
[Cu(H2O)Cl3](aq) + Cl(aq)
K1 = 3.1 mol1 dm3
[CuCl4]2(aq) + H2O(l)
[Cu(H2O)4]2+(aq) + 4Cl(aq)
[CuCl4]2(aq) + 4H2O(l)
K1 = 6.3102 mol1 dm3
K2 = 40 mol1 dm3
Kst = K1  K2  K3  K4 = 4.2  105 mol4 dm12
165
K1 > K2 > K3 > K4
Reasons :
1. Statistical effect
On successive displacement, less water
ligands are available to be displaced.
166
K1 > K2 > K3 > K4
Reasons :
2. Charge effect
On successive displacement, the Cl
experiences more repulsion from the
complex
[Cu(H2O)4]2+
[Cu(H2O)Cl3]
167
Cl
Cl
attraction
repulsion
Colours of some copper(II) complexes
Formula of copper(II) complex
Colour of the
complex
[Cu(H2O)4]2+
Pale blue
[CuCl4]2–
Yellow
[Cu(NH3)4]2+
Deep blue
[Cu(H2NCH2CH2NH2)]2+
[Cu(EDTA)]2–
Violet
Sky blue
The displacement of ligands are usually
accompanied with easily observable colour changes
168
Coloured Ions
The colours of many gemstones are due to the
presence of small quantities of d-block metal ions
169
Coloured Ions
•
Most of the d-block metals
 form coloured compounds
 due to the presence of the
incompletely filled d orbitals in the
d-block metal ions
Which aqueous transition metal ion(s) is/are
not coloured ?
3d10 : Zn2+, Cu+;
170
3d0 : Sc3+, Ti4+
Colours of some d-block metal ions in aqueous solutions
Number of unpaired
electrons in 3d
orbitals
0
1
171
d-Block metal
ion
Colour in
aqueous solution
Sc3+
Colourless
Ti4+
Colourless
Zn2+
Colourless
Cu+
Colourless
Ti3+
Purple
V4+
Blue
Cu2+
Blue
Colours of some d-block metal ions in aqueous solutions
Number of unpaired
electrons in 3d
orbitals
2
3
172
d-Block metal
Colour in
ion
aqueous solution
V3+
Green
Ni2+
Green
V2+
Violet
Cr3+
Green
Co2+
Pink
Colours of some d-block metal ions in aqueous solutions
Number of unpaired
electrons in 3d
orbitals
4
5
173
d-Block metal
Colour in
ion
aqueous solution
Cr2+
Blue
Mn3+
Violet
Fe2+
Green
Mn2+
Very pale pink
Fe3+
Yellow
Co2+(aq)
Zn2+(aq)
Fe3+(aq)
Colours of some d-block metal ions in aqueous solutions
174
Mn2+(aq)
Fe2+(aq)
Cu2+(aq)
Colours of some d-block metal ions in aqueous solutions
175
A substance absorbs visible light of a certain
wavelength
 reflects or transmits visible light of
other wavelengths (complimentary colour)
 appears coloured
Light
Light reflected or
absorbed
transmitted
[Cu(H2O)4]2+(aq)
Yellow
Blue
[CuCl4]2(aq)
Blue
Yellow
Coloured ion
176
Complimentary colour
chart
Blue light absorbed
Appears yellow
Violet
Blue
Cyan
Yellow light absorbed
Appears blue
177
Green
Magenta
Red
Yellow
Greenish yellow
Coloured Ions
•
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
178
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
179
distributes along z axis
Interact more strongly with
the orbitals of ligands
distributes along x and y axes
d z 2 , d x2 y 2
eg
dxy , dxz , dyz
The splitting of the degenerate 3d orbitals of
a d-block metal ion in an octahedral complex
180
t 2g
dx
181
2
y 2
Higher energy eg
Criterion for d-d transition : presence of unpaired d electrons in the dblock 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
182
H2O as ligand

Cu2+

3d9 : d-d transition is possible
183
Yellow light absorbed,
appears blue

H2O as ligand
*Cu2+

3d9 : d-d transition is possible
184
Fe2+

3d6 : d-d transition is possible
185
Magenta light absorbed,
appears green

*Fe2+

3d6 : d-d transition is possible
186

Zn2+

3d10 : d-d transition NOT possible
187
Sc3+
3d0 : d-d transition NOT possible
188
E
E depends on
[Fe(H2O)6]2+ green,
[Fe(H2O)6]3+ yellow
1. the nature and charge of metal ion
2. the nature of ligand
189
[Cu(H2O)4]2+ blue,
[CuCl4]2
yellow
Coloured Ions
Why does Na+(aq) appear colourless ?
3d0 : d-d transition is NOT possible
2p  3s transition involves absorption
of radiation in the UV region.
190
Catalytic Properties of Transition
Metals and their Compounds
•
The d-block metals and their compounds
 important catalysts in industry and
biological systems
191
The use of some d-block metals and their compounds as
catalysts in industry
d-Block
metal
Catalyst
V
V2O5 or
vanadate(V) (VO3–)
Reaction catalyzed
Contact process
2SO2(g) + O2 (g)
2SO3(g)
Haber process
Fe
192
Fe
N2(g) + 3H2(g)
2NH3(g)
The use of some d-block metals and their compounds as
catalysts in industry
d-Block
metal
Ni
Pt
193
Catalyst
Reaction catalyzed
Ni
Hardening of vegetable oil
(Manufacture of margarine)
RCH = CH2 + H2  RCH2CH3
Pt
Catalytic oxidation of ammonia
(Manufacture of nitric(V) acid)
4NH3(g) + 5O2(g)  4NO(g) + 6H2O(l)
Catalytic Properties of Transition
Metals and their Compounds
•
The d-block metals and their compounds
exert their catalytic actions in either
 heterogeneous catalysis
 homogeneous catalysis
194
Catalytic Properties of Transition
Metals and their Compounds
•
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
195
1. Heterogeneous Catalysis
•
The catalyst and reactants
 exist in different states
•
The most common heterogeneous
catalysts
 finely divided solids for gaseous
reactions
196
1. Heterogeneous Catalysis
A heterogeneous catalyst provides a
suitable reaction surface for the reactants
to come close together and react.
197
1. Heterogeneous Catalysis
•
Example:
The synthesis of gaseous ammonia from
nitrogen and hydrogen (i.e. Haber
process)
N2(g) + 3H2(g)
198
2NH3(g)
1. Heterogeneous Catalysis
•
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
199
1. Heterogeneous Catalysis
•
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
200
1. Heterogeneous Catalysis
•
In the presence of iron as catalyst
 the reaction proceeds much faster
 provides an alternative reaction
pathway of lower activation energy
201
1. Heterogeneous Catalysis
•
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
202
1. Heterogeneous Catalysis
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
203
1. Heterogeneous Catalysis
2. The iron metal
 many 3d electrons and low-lying
vacant 3d orbitals
 form bonds with the reactant
molecules
 adsorb them on its surface
204
 weakens the bonds present in the
reactant molecules
1. Heterogeneous Catalysis
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
205
The catalytic mechanism of the formation of
gaseous ammonia from nitrogen and hydrogen
206
The catalytic mechanism of the formation of
gaseous ammonia from nitrogen and hydrogen
207
The catalytic mechanism of the formation of
gaseous ammonia from nitrogen and hydrogen
208
The catalytic mechanism of the formation of
gaseous ammonia from nitrogen and hydrogen
209
The catalytic mechanism of the formation of
gaseous ammonia from nitrogen and hydrogen
210
43.3 Characteristic Properties of the d-Block Elements and their compound
(SB p.162)
1. Heterogeneous Catalysis
•
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
211
Energy profiles of the reaction of nitrogen and hydrogen to
form gaseous ammonia in the presence and absence of a
heterogeneous catalyst
212
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
213
2. Homogeneous Catalysis
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 (S2O82–) and iodide ions (I–)
214
2. Homogeneous Catalysis
•
Peroxodisulphate(VI) ions
 oxidize iodide ions to iodine in an
aqueous solution
 themselves being reduced to
sulphate(VI) ions
S2O82–(aq) + 2I–(aq)
2SO42–(aq) + I2 (aq)
o
cell
E
215
 1.51V
2. Homogeneous Catalysis
•
•
The reaction is very slow due to strong
repulsion between like charges.
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)
216
= +0.23 V
2. Homogeneous Catalysis
•
Iron(II) ions
 subsequently oxidized by
peroxodisulphate(VI) ion
 the original iron(III) ions are
regenerated
2Fe2+(aq) + S2O82–(aq)
2Fe3+(aq) + 2SO42–(aq)
217
= +1.28 V
2. Homogeneous Catalysis
•
The overall reaction:
2I–(aq) + 2Fe3+(aq)
I2(aq) + 2Fe2+(aq)
+)
= +0.23 V
2Fe2+(aq) + S2O82–(aq)
2Fe3+(aq) + 2SO42–(aq)
= +1.28 V
S2O82–(aq) + 2I–(aq)
2SO42–(aq) + I2(aq)
= +1.51 V
Feasible reaction
218
43.3 Characteristic Properties of the d-Block Elements and their compound
(SB p.164)
2. Homogeneous Catalysis
•
Iron(III) ions
 catalyze the reaction
 acting as an intermediate for the
transfer of electrons between
peroxodisulphate(VI) ions and iodide
ions
219
2. Homogeneous Catalysis
•
Peroxodisulphate(VI) ions
 oxidize Fe2+ to Fe3+
•
Iodide ions
 reduce Fe3+ to Fe2+
220
The End
221
Check Point 43-3E
222
Energy profiles
for the oxidation
of iodide ions by
peroxodisulphate
(VI) ions in the
presence and
absence of a
homogeneous
catalyst
Besides iron(III) ions, iron(II) ions can also catalyze
the reaction between peroxodisulphate(VI) ions and
iodide ions. Why?
Answer
223
Iron(II) ions catalyze the reaction by reacting with the
peroxodisulphate(VI) ions first.
2Fe2+(aq) + S2O82–(aq)
2Fe3+(aq) + 2SO42–(aq)
The iron(III) ions formed then oxidize the iodide ions.
2Fe3+(aq) + 2I–(aq)
2Fe2+(aq) + I2(aq)
In this way, the reaction between peroxodisulphate(VI) ions and
iodide ions is catalyzed.
Back
224
Which of the following redox systems might catalyze the
oxidation of iodide ions by peroxodisulphate(VI) ions in
an aqueous solution?
Cr2O72–(aq) + 14H+(aq) + 6e–
2Cr3+(aq) + 7H2O(l)
Answer
= +1.33 V
MnO4–(aq) + 8H+(aq) + 5e–
Mn2+(aq) + 4H2O(l)
= +1.52 V
Sn4+(aq) + 2e–
= +0.15 V
Sn2+(aq)
(Given: S2O82–(aq) + 2e–
225
I2(aq) + 2e–
2SO42–(aq)
2I–(aq)
= +2.01 V
= +0.54 V)
Those redox systems with
greater than +0.54 V and smaller than
+2.01 V are able to catalyze the oxidation of iodide ions by
peroxodisulphate(VI) ions in an aqueous solution. Therefore, the
following two redox systems are able to catalyze the reaction.
Cr2O72–(aq) + 14H+(aq) + 6e–
2Cr3+(aq) + 7H2O(l)
MnO4–(aq) + 8H+(aq) + 5e–
Mn2+(aq) + 4H2O(l)
Back
226