Download TRANSITION METALS

TRANSITION METALS
The transition metals occupy the central block of the Periodic Table.
A transition element is an element having an incomplete d-shell in the
element or in one of its other oxidation states.
This means that not all d-block elements are classed as transition metals, for example
zinc. Zinc has two oxidation states: 0 ([Ar]4s2 3d10) and +2 ([Ar]3d10). In both of these
oxidation states it has a complete d-shell.
When a transition element in the first row loses electrons, electrons in the
4s level are lost before those in the 3d level.
In the first transition series, there are ten elements (Sc-Zn), corresponding to the filling
of the 3d sub-shell. The electronic configurations of these elements and some of their
common ions are given in the following table.
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
4s2 3d1
4s2 3d2
4s2 3d3
4s1 3d5
4s2 3d5
4s2 3d6
4s2 3d7
4s2 3d8
4s1 3d10
4s2 3d10
Ti2+
V2+
Cr2+
Mn2+
Fe2+
Co2+
Ni2+
Cu2+
Zn2+
Sc3+
Ti3+
V3+
Cr3+
Mn3+
Fe3+
Co3+
3d2
3d3
3d4
3d5
3d6
3d7
3d8
3d9
3d10
3d0
3d1
3d2
3d3
3d4
3d5
3d6
Ti4+
V4+
3d0
3d1
Cu+
3d10
non-transitional
Higher oxidation states usually occur in the form of oxo anions, for example
manganate(VII) MnO4-.
General Properties
1. Appearance
The d-block elements are all metals, and with the exception of copper and gold, are all
silvery-white in colour.
2. Atomic radius
Since the electrons are entering the inner 3d-subshell and the nuclear charge is
increasing, there is initially a contraction in the atomic radius as the d-block is crossed.
Eventually there is expansion, caused by increased screening of the outer electrons.
Since there is little variation in atomic radius from vanadium to copper, these metals
readily alloy with one another.
TOPIC 13.23: THE TRANSITION METALS
1
3. Density
The transition elements have small atomic radii and atomic volumes and high nuclear
masses. As a result, their densities are high in comparison to the preceding s-block
elements.
8
6
density
g.cm-3
4
2
0
Sc
Ti
V
Cr
Mn Fe
Co
Ni
Cu
Zn
4. First Ionisation Energy
The similarity of the d-block elements is exemplified by their similar ionisation energies.
The slow increase in I1 across the d-block is caused by the increasing nuclear charge.
1st
ionisation
energy
kJ.mol-1
800
400
0
Sc
Ti
V
Cr
Mn Fe
Co
Ni
Cu
Zn
5. Melting Point
Owing to the large number of 3d and 4s electrons involved in the metallic bonding, these
metals have high melting points. The exception is zinc, where only the 4s electrons are
involved in bonding.
All these metals have high electrical conductivities.
TOPIC 13.23: THE TRANSITION METALS
2
3d
Sc
[Ar]
Ti
[Ar]
V
[Ar]
Cr
[Ar]
Mn
[Ar]
Fe
[Ar]
Co
[Ar]
Ni
[Ar]
Cu
[Ar]
Zn
[Ar]
4s
As the outer 4s shell for all
these atoms is occupied,
they present a more or less
similar face to the world.
Similarities exist across the
period as well as down the
group.
Cr and Cu have anomalous
configurations, which arise
because of the enhanced
stability of half-filled dorbitals.
This stability also explains
why Fe2+ ([Ar]3d6) is easily
oxidised to Fe3+ ([Ar]3d5)
and Mn2+ ([Ar]3d5) is not
easily oxidised to Mn3+
([Ar]3d4)
Characteristic properties of transition elements are:

variable oxidation states

formation of coloured ions

formation of complexes

catalytic activity
TOPIC 13.23: THE TRANSITION METALS
3
Formation of Complexes
Complex compounds consist of a central metal atom surrounded by ligands.
A ligand is an electron pair donor.
A ligand may be an atom, an ion or a molecule.
N.B. Ligands, nucleophiles and Lewis bases are all defined as electron pair donors.
Since the ligand donates both of the electrons it shares with the central metal ion, a
coordinate bond is formed.
The co-ordination number is the number of atoms bonded to the metal ion.
In a complex, the co-ordination number of the metal differs from its oxidation state.
Monodentate Ligands
Some ligands have only one atom which can donate a pair of electrons, they therefore
form only one bond with the metal ion. These are described as monodentate ligands;
examples of monodentate ligands are:
CN-
NH3
OH-
H2O
F-
Cl-
Br-
For a complex which is formed from a metal ion and monodentate ligands only, the
number of ligands is equal to the co-ordination number.
e.g.
hexaaquairon(II) ion, [Fe(H2O)6]2+
H
H
2+
..
O
H
:O
O:
H
H
H
Fe
O:
:O
H
H
H
no. of ligands = 6
co-ordination no. = 6
H
..
shape: octahedral
oxidation state of iron = +2
O
H
H
The Fe-O and O-H bonds within this complex ion are covalent bonds.
TOPIC 13.23: THE TRANSITION METALS
4
Bidentate Ligands
Some ligands contain two atoms which can donate a pair of electrons, they are therefore
able to form two bonds with the metal ion. These are described as bidentate ligands;
examples of bidentate ligands are:
O:
NH2
..
-
O
CH2
C
CH2
C
..
..N
O
-
O:
NH2
1,2-diaminoethane (en)
e.g
..N
ethanedioate ion
bipyridyl (bipy)
[Ni(en)3]2+
CH2
NH2
..
CH2
2+
CH2
shape: octahedral
:NH2
H2N:
CH2
Ni
H2N:
co-ordination no. = 6
:NH2
..
oxidation state of nickel = +2
CH2
NH2
no. of ligands = 3
CH2
The number of ligands is no longer equal to the co-ordination number.
NB where there are 2 or more bidentate ligands in a complex it is possible to make two
non-superimposable mirror images and the molecules are said to be chiral. Chiral
molecules rotate the plane of plane-polarised light in opposite directions.
Optical activity will be covered in more detail in topic shortly in organic chemistry.
TOPIC 13.23: THE TRANSITION METALS
5
Multidentate Ligands
Some ligands contain many atoms which can donate a pair of electrons, they are
therefore able to form many bonds with the metal ion. These are described as
multidentate ligands. The most important multidentate ligand is EDTA4(ethylenediaminetetracetate), which has six donor atoms. It is the anion of
bis[di(carboxymethyl)amino]ethane:
..
..
-
-OOCCH
2
..
..
NCH2CH2N
CH2COO
..
..
CH2COO-
OOCCH2
EDTA forms 1:1 octahedral complexes with M2+ ions, for example:
[Cu(H2O)6]2+ + EDTA4-
[Cu(EDTA)]2- + 6H2O
The Chelate effect
In the above example we saw a single hexadentate ligand replace 6 monodentate
ligands. The enthalpy change for the above reaction is not expected to be very
significant since 6 metal-oxygen co-ordinate bonds are being broken and 4 metaloxygen and 2 metal-nitrogen co-ordinate bonds are being formed. Such a reaction,
however, occurs with a significant gain in entropy as 2 particles become seven. For this
reason the formation of a complex with a multidentate ligand over a monodentate ligand
is favoured. This is called the chelate effect.
TOPIC 13.23: THE TRANSITION METALS
6
Haemoglobin
The red pigment haem, which
is a complex of iron(II), is found
in blood. The iron is bonded to
the four nitrogen atoms of a
planar porphyrin molecule. The
porphyrin
molecule
is
a
quadridentate ligand.
N
N
Fe
In haemoglobin, the iron has a
coordination number of 6, with
octahedral
geometry.
In
addition to its four bonds to the
porphyrin molecule, iron bonds
to a fifth nitrogen atom, from
the protein globin, and to either
an
oxygen
molecule
(in
oxygenated blood) or a water
molecule (in deoxygenated
blood). Changing the ligand
from O2 to H2O changes the
colour of the complex from
bright red to red-purple.
N
N
Strong field ligands, such as CO and CN- are able to bond more strongly to haem than
oxygen can; they can therefore prevent the blood from carrying oxygen. This explains
the high toxicity of carbon monoxide and cyanides.
Shapes of Complex Ions
Co-ordination number 2
Silver ions and copper(I) ions, both of which have a d 10 configuration, are unusual in
forming linear complexes by bonding to only two ligands. The bond angle is 180 o.
[Ag(NH3)2]+ This complex ion, which is formed when either AgCl or Ag 2O dissolves in
excess aqueous ammonia, is colourless because of the d 10 configuration of silver(I). It is
the ion present in Tollen’s reagent. Tollen’s reagent is used to distinguish aldehydes
from ketones: when aldehydes are heated with Tollen’s reagent, the complex is reduced
to metallic silver, which is deposited on the test tube walls as a mirror. There is no
reaction with ketones.
RCHO + 2[Ag(NH3)2]+ + 3OHTOPIC 13.23: THE TRANSITION METALS
RCOO- + 2Ag +4NH3 +2H2O
7
Co-ordination number 4
Where a ligand is large and negatively charged, such as Cl -, Br-, I-, there is not usually
enough space to fit six ligands around the central metal atom, and four ligands bond
instead. Complexes with a co-ordination number of four are usually tetrahedral in shape.
Examples:
2-
2-
Cl
Cl
Cu
Co
Co
Cl
Cl
-
Cl
Cl
Cl
Cl
Cl
[CuCl4]2yellow-green
[CoCl4]2blue
Fe
Cl
Cl
Cl
[FeCl4] colourless
Another possible shape for four ligands around a central metal ion is square planar.
Although generally rare, this geometry occurs in some nickel complexes, such as
[Ni(CN)4]2-, and is common among complexes of platinum. An important example is
cisplatin, [PtCl2(NH3)2].
Cl
Cl
Pt
H3N
NH3
Cisplatin (so called because the two Cl groups are on the
same side of the molecule) is an anti-cancer drug. It is
used in chemotherapy for certain types of cancer, such as
testicular cancer, and has had a major success rate.
The other geometric isomer called transplatin (Cl on
opposite sides) has no such properties.
NB cis and trans isomerism is a type E/Z isomerism which you have already met.
Co-ordination number 6
The most common shape for complex ions is octahedral; it is found with a wide variety
of ligands and metals. For example:
[Cu(H2O)6]2+
[Ni(NH3)6]2+
[Fe(CN)6]3[Cr(OH)6]3- [PtCl6]2- [Co(NO2)6]3All the transition metals of the first row transition series form hexaaqua ions which are
octahedral.
TOPIC 13.23: THE TRANSITION METALS
8
Formation of Coloured Ions
Crystal Field Theory
In a metal ion such as Ti3+, which contains a single d-electron, all five d-orbitals are
equivalent, and the electron has an equal probability of being in any one of them.
However, if a metal ion, Mn+, lies at the centre of an octahedral set of charges, such as
when it is surrounded by six identical ligands (L), the d-orbitals are no longer equivalent.
A ligand is an electron pair donor.
L
L
Mn+
L
L
L
L
This is because the dz2 and dx2-y2
(degenerate) orbitals have lobes which are
heavily concentrated in the vicinity of the
charges, whereas the dxy, dyz and dxz orbitals
have lobes which project between the
charges. Since the electron experiences less
repulsion in these latter three (degenerate)
orbitals, they will have lower energy than the
former two orbitals. The diagram below
shows the effect of the octahedral
environment of six ligands.
}
0.6o
o
{
5 degenerate d-orbitals
in the absence of ligands
0.4o
}
d-orbital splitting in the presence of ligands
If the equivalent reasoning is applied to a metal ion, Mn+, which lies at the centre of a
tetrahedral set of charges, such as when it is surrounded by four identical ligands, it can
be seen that the degenerate orbitals dxy, dyz and dxz have higher energy than the
degenerate orbitals dz2 and dx2-y2.
TOPIC 13.23: THE TRANSITION METALS
9
}
0.4t
{
t
5 degenerate d-orbitals
in the absence of ligands
0.6t
}
d-orbital splitting in the presence of ligands


is usually greater thant. The energy differences o and t correspond to frequencies
in the visible part of the electromagnetic spectrum (E = hhc). Therefore, if an
electron moves from the lower to the higher energy level, visible light is absorbed and
the substance is coloured. In the above expression, h = Plank’s constant.
The factors which affect the size of the d level splitting and therefore the colour
are:
 the nature of the metal
 the oxidation state of the metal
 the nature of the ligand
 the co-ordination number (i.e. the number of ligands)
For example:
[CoIII(NH3)6]3+
yellow
octahedral
[CoII(H2O)6]2+
pink
octahedral
[CoIICl4]2blue
tetrahedral
Ligands are said to possess a ligand field; the greater the ligand field, the greater the
splitting of the d levels. Ligands can thus be arranged into the Spectrochemical Series:
CNNH3
most splitting
[CrCl6]3[Cr(H2O)6]3+
[Cr(NH3)6]3+
[Cr(CN)6]3-
H2O
OH-
F-
max /nm
E /kJ.mol-1
733
575
461
381
163
208
259
314
TOPIC 13.23: THE TRANSITION METALS
10
Cl-
Brleast splitting
violet
ruby
purple
green
Change of metal
With the same ligand, different metals in the same oxidation state have different colours:
max /nm
E /kJ.mol-1
[Fe(H2O)6]3+
[Cr(H2O)6]3+
[Mn(H2O)6]3+
730
575
476
164
208
251
pale violet
ruby
purple
[Fe(H2O)6]2+
[Cr(H2O)6]2+
[Co(H2O)6]2+
962
719
540
124
166
221
pale blue-green
blue
pink
Change of oxidation state
The same ligand with the same metal, but in different oxidation states, produces
different colours:
max /nm
E /kJ.mol-1
[Fe(H2O)6]2+
[Fe(H2O)6]3+
962
730
124
164
pale blue-green
pale violet
[Cr(H2O)6]2+
[Cr(H2O)6]3+
719
575
166
208
blue
ruby
Change of ligand
A different ligand with the same metal in the same oxidation state gives rise to different
colours.
max /nm
E /kJ.mol-1
[CrCl6]3[Cr(H2O)6]3+
[Cr(NH3)6]3+
[Cr(CN)6]3-
733
575
461
381
163
208
259
314
violet
ruby
purple
green
Change of co-ordination number
Small ligands are able to form octahedral complexes, in which six ligands surround the
central metal ion. With larger ligands, there is not enough room to accommodate six
around the central metal ion. Instead, a tetrahedral complex is formed, in which four
ligands surround the central metal ion.
[Co(H2O)6]2+
[CoCl4]2-
max /nm
E /kJ.mol-1
540
690
221
173
TOPIC 13.23: THE TRANSITION METALS
11
pink
blue
Ultraviolet & visible spectrophotometry
The intensity of colours given by a solution of a transition metal ion can be used to find
the concentration of the ion.
In an ultraviolet/ visible spectrophotometer, the amount of light absorbed at a particular
wavelength is proportional to the concentration of the species which is absorbing.
The Beer-Lambert law:
A= cl
A is the absorbance, c is the concentration in mol.dm-3, l is the path length in
cm and  is the molar absorption coefficient
, the molar absorption coefficient, is a measure of the intensity of the colour; the greater
the value of , the more intense the colour. Some transition metal ions, for example
[Fe(H2O)6]2+ [Fe(H2O)6]3+ and [Mn(H2O)6]2+, have very low values of , and therefore
accurate measurements of concentration by spectrophotometry are not possible. In such
cases, changing the ligand to one which gives a much more intense colour makes
accurate measurement easy and does not alter the concentration of the metal ion. In the
case of iron, the ligand bipyridyl (bipy) can be used to intensify the colour of Fe(II) and
the ligand thiocyanate to intensify the colour of Fe(III).
[Fe(H2O)6]2+ + 3 bipy
pale blue-green
[Fe(bipy)3]2+ + 6H2O
deep red
[Fe(H2O)6]3+ + SCNpale violet
[Fe(H2O)5(SCN)]2+ + H2O
blood-red
N
..
N
..
bipyridyl
Measurement of the colour could be carried out in a colorimeter. A coloured filter,
corresponding to the wavelength of maximum absorbance, is placed in the machine,
and allows absorbance to be measured over a narrow band of wavelengths. A
spectrophotometer is a more sophisticated machine. It scans through the visible and
ultraviolet regions, measuring the absorbance at every wavelength.
TOPIC 13.23: THE TRANSITION METALS
12
Variable Oxidation States
Except for scandium, which has exclusively an oxidation state of +3, the first transition
series elements show an oxidation state of +2 (when both the 4s-electrons have been
lost) and higher oxidation states, when 3d-electrons are involved.
As the period is traversed, there is a trend towards decreased stability of the very high
oxidation states and also the +2 oxidation state becomes more stable relative to the +3
oxidation state. In general, a +3 aqueous ion will oxidise the +2 aqueous ion of the
element to the left of it. The exception is that Mn3+(aq) oxidises Fe2+(aq).
The variable oxidation states displayed by transition metals are primarily due to the fact
that successive ionisation energies of the atom increase gradually. For s and p block
metals, there is a point at which there is a sharp increase in the values of successive
ionisation energies.
12000
10000
8000
ionisation
6000
energy
kJ.mol-1
I4
I3
I2
4000
I1
2000
0
Al
V
TOPIC 13.23: THE TRANSITION METALS
Cr
Fe
13
Zn
The following oxidation states are known. The most stable oxidation state is circled; the
most common oxidation states are in red.
Sc
3
Ti
V
Cr
Mn
7
Fe
Co
Ni
6
6
6
5
5
5
4
4
4
4
4
4
3
3
3
3
3
3
2
2
2
2
2
2
2
Cu
Zn
2
2
1
In general, oxidation states lower than the most stable one are reducing agents, while
those which are higher are oxidising agents.
Oxidation states of vanadium
Ammonium vanadate(V), NH4VO3, is a white solid. When added to dilute hydrochloric
acid, the orange-yellow colour of the dioxovanadate(V) ion, VO2+, is formed.
VO3- + 2H+
VO2+ + H2O
There is no change in the oxidation state of vanadium.
When this solution is shaken with zinc amalgam, colour changes take place over several
minutes as vanadium is reduced through a number of oxidation states.
VO2+(aq) + 2H+(aq) + e-
VO2+(aq) + H2O
blue
VO2+(aq) + 2H+(aq) + eV3+(aq) + e-
V3+(aq) + H2O
green
V2+(aq)
violet
The final stage of the reduction, from vanadium(III) to vanadium(II), is best carried out
under an atmosphere of nitrogen, since vanadium(II) ions are oxidized by oxygen in air.
TOPIC 13.23: THE TRANSITION METALS
14
Oxidation states of manganese
The highest oxidation state of manganese is +7, which is found in manganate(VII) ions,
MnO4-. This ion is an intense purple colour.
When this solution is acidified and shaken with a reducing agent, it is reduced directly to
manganese(II) ions. [Mn(H2O)6]2+ ions are very pale pink and appear almost colourless
in aqueous solution.
MnO4-(aq) + 8H+(aq) + 5e-
Mn2+(aq) + 4H2O(l)
colourless
Manganate (VII) is commonly used in redox titrations to determine the concentration of
iron (II). The intense purple colour disappears as the Mn 2+ ion is formed as per the
equation above and so acts as its own indicator. The choice of acid used in redox
titrations using MnO4- is very important. Consider the redox potentials below;
Eo (V)
+1.51
+1.36
+0.81
+0.77
+0.17
Reduction half equation
MnO4- +8H+ + 5e-  Mn2+ + 4H2O
Cl2 + 2e-  2ClNO3- + 2H+ + e-  NO2 + H2O
Fe3+ +e-  Fe2+
SO42- + 4H+ + 2e-  H2SO3 + H2O
If nitric acid is used to acidify the manganate (VII) solution, nitrate ions will oxidise iron
(II) to iron (III) before any manganate (VII) is added. The titre in this titration will
therefore be very much lower than it should be.
Fe2+ + NO3- + 2H+  Fe3+ + NO2 + H2O
(Eo = +0.04V)
If hydrochloric acid is used to acidify the manganate (VII) solution, chloride ions will be
oxidized to chlorine by the manganite (VII) ions as well as the iron (II). The titre in this
titration will therefore be very much higher than it should be.
10Cl- + 2MnO4- +16H+  2Mn2+ + 8H2O + 5Cl2
(Eo = +0.15V)
Only sulfuric acid can be used for this titration.
Under alkaline conditions, manganate(VII) ions are reduced to manganese(IV) oxide
which appears as a brown precipitate.
MnO4-(aq) + 2H2O(l) + 3e-
TOPIC 13.23: THE TRANSITION METALS
MnO2(s) + 4OH-(aq)
15
Oxidation in alkaline solution
Under alkaline conditions, transition metals in low oxidation states can be oxidised to
higher oxidation states. For example:
If sodium hydroxide solution is added to a solution of manganese(II) ions, the initial offwhite precipitate of manganese(II) hydroxide (Mn(OH)2) gradually darkens on standing
as it is oxidised by air to manganese(IV) oxide (MnO2).
Similarly, if sodium hydroxide solution is added to a solution of iron(II) ions, the initial
dark green precipitate of iron(II) hydroxide (Fe(OH)2) goes brown fairly quickly on
standing as it is oxidised by air to iron(III) hydroxide (Fe(OH)3).
Oxidation is the loss of electrons. It is easier to lose an electron from a negative ion than
from a neutral molecule; it is easier to lose an electron from a neutral molecule than
from a positive ion. The metal-containing species present at different pH’s are:
[M(H2O)6]2+
in acid solution
[M(H2O)4(OH)2]
neutral
[M(H2O)2(OH4)]2in alkaline solution
Clearly, oxidation will take place most easily under alkaline conditions.
Alkaline oxidation:
 add an excess of sodium hydroxide solution
 add an oxidising agent, such as hydrogen peroxide
An example is the oxidation of chromium(III). When excess NaOH is added to a solution
of a chromium(III) salt, a dark green solution containing [Cr(OH)6]3- is formed. Treatment
of this solution with hydrogen peroxide gives a yellow solution containing the
chromate(VI) ion, CrO42-.
2[Cr(OH)6]3- + 3H2O2
2 CrO42- + 2OH- + 8H2O
Another example is the reaction which occurs when concentrated ammonia solution is
added to a pink solution of cobalt(II) ions until in excess. The initial green-blue
precipitate of cobalt(II) hydroxide, Co(OH)2, re-dissolves to give a pale brown solution
containing the hexaammine cobalt(II) ion. In the presence of air, the solution rapidly
darkens and becomes dark brown as oxidation to the hexaammine cobalt(III) ion takes
place. Although the hexaammine cobalt(III) ion is yellow, the mixture is dark because
other species are present.
[Co(H2O)6]2+
pink
[Co(NH3)6]2+
pale brown
[Co(H2O)4(OH)2]
green-blue
TOPIC 13.23: THE TRANSITION METALS
16
[Co(NH3)6]3+
yellow
Catalysis
A catalyst is a substance which alters the rate of a chemical reaction
but is not used up in the reaction.
Most catalysts are positive catalysts – they speed up chemical reactions. However,
negative catalysts are sometimes used to retard undesirable reactions. For example,
phosphoric acid is added to hydrogen peroxide to slow down the rate of its
decomposition in light.
A catalyst works by providing an alternative reaction path
with a lower activation energy.
Catalysts are usually divided into two classes:
In heterogeneous catalysis, the catalyst and reactants are in different phases.
In homogeneous catalysis, the catalyst and reactants are in the same phase.



A catalyst increases the rate of attainment of equilibrium but does not alter the
position of equilibrium.
A substance which reduces the efficiency of a catalyst is called an inhibitor or a
catalyst poison. Catalyst poisoning occurs with heterogeneous catalysts when,
for example, sulphur compounds block the active sites; the sulphur is converted
to inactive sulphides on the catalyst surface. Catalytic converters in cars are
poisoned by any lead which is present in the petrol.
Catalysts are important in improving the economics of industrial processes, since they
allow products to be made in a shorter time. Transition metals are effective catalysts
owing to their ability to exhibit different oxidation states and are widely used.
Examples of transition elements as catalysts:
Element
Catalyst
Reaction
Ti
V
Cr
Mn
Fe
Fe
Co
Ni
Cu
Pt
TiCl4 / Al(C2H5)3
V2O5
Cr2O3
MnO2
Fe
Fe2O3
Co
Ni
CuO
Pt
Ziegler-Natta polymerisation of alkenes
Contact process
CO + 2H2
CH3OH
Decomposition of H2O2
Haber process
Bosch process
Oxo process: RCH=CH2
RCH2CH2CHO
Catalytic hydrogenation of alkenes
Deacon process: 4HCl + O2
2H2O + Cl2
Ostwald process; catalytic converters
TOPIC 13.23: THE TRANSITION METALS
17
Heterogeneous Catalysis
Heterogeneous catalysis typically involves gaseous reactants combining in the presence
of a solid catalyst. Most industrial processes use heterogeneous catalysts.
Heterogeneous catalysis can be regarded as involving three separate stages:
1. Adsorption
In a simple gas phase reaction: X(g) + Y(g)
Z(g)
at least one of the reactants is adsorbed onto active sites on the surface of the catalyst
by the formation of weak chemical bonds. This can result in an increase in reaction rate
in a number of ways:
 X alone is adsorbed onto the surface of the catalyst, where it becomes effectively
more concentrated. Molecules of X are therefore more likely to undergo collisions
with Y than they would be purely in the gas phase - the frequency of collisions
is increased
 existing chemical bonds may be weakened, making them easier to break - Ea is
lowered
 X and Y are both adsorbed onto the surface of the catalyst, where they are in
contact with each other at concentrations far greater than in the surrounding gas the frequency of collisions is increased
 the correct spacing of the active sites can ensure that molecules are held in the
right geometry to react – Ea is lowered
2. Reaction
The chemical reaction takes place on the surface of the catalyst.
3. Desorption
The products of the reaction must be desorbed from the active sites to allow further
molecules of reactant to be adsorbed.
Preparing the catalyst
Since a chemical reaction takes place only on the surface of the catalyst, increasing the
surface area improves the efficiency of the catalyst. This is usually achieved by
spreading the catalyst very thinly on an inert support medium. For example, in catalytic
converters in cars, the very expensive mixed catalyst of platinum and rhodium is spread
thinly onto a cheap ceramic support. The surface to mass ratio of the catalyst is greatly
increased. Loss of catalyst from the surface of the support is a problem and can greatly
increase the overall cost of the catalyst.
Strength of Adsorption
The strength of the adsorption onto the surface of the catalyst determines the
effectiveness of the catalyst.
If the adsorption is too weak, the rate of a reaction may not be greatly enhanced
because:
 existing bonds are not weakened sufficiently
 molecules may not be held in the right geometry
 the concentration of reactant(s) on the catalyst surface is not increased enough
TOPIC 13.23: THE TRANSITION METALS
18
Too weak adsorption tends to occur with metals near the end of a transition series, such
as silver.
Occasionally, weak adsorption can be advantageous, when one reaction is favoured
above other competing and unwanted processes. An example of this is the silvercatalysed oxidation of ethene to epoxyethane.
Reactant molecules must be free to move around on the surface of the catalyst in order
to meet other reactant molecules and react. Products must be released fast enough to
increase the overall reaction rate. If the adsorption is too strong,
 reactant molecules cannot move freely over the catalyst surface
 product molecules may not be desorbed from the surface of the catalyst quickly
enough, if at all, blocking the active sites. The catalyst becomes poisoned by
product molecules, and the reaction rate is not increased.
Too strong adsorption tends to occur with metals at the beginning of a transition series,
such as tungsten.
Metals in the middle of a transition series, such as nickel and platinum, achieve a
balance between the two effects and are good catalysts.
Volcano curves show the change in catalytic activity across a transition series.
x105
x
x
x104
x
Relative
reaction rate
x103
x
x102
x
x
x
x
x10
x1
x
x
Nb
Cr
Mo
TOPIC 13.23: THE TRANSITION METALS
x
Mn
Tc
Fe
Ru
19
Co
Rh
Ni
Pd
x
Cu
Specific Surface Catalysis
Different distances between active sites in different catalysts and their different affinities
for reactant molecules result in heterogeneous catalysts being specific. Different
catalysts can bring about different reactions with the same reactant(s). For example:
In the presence of a copper catalyst, ethanol is
oxidised to ethanal. The adsorption of the O-H group
weakens the hydrogen-oxygen bond, favouring
dehydrogenation.
CH3.CH2
O
H
Cu
CH3
H
H
C
H
O
Al2O3
In the presence of an aluminium oxide catalyst,
ethanol reacts to form ethene. The adsorption of the
C-O group weakens the carbon-oxygen bond,
favouring dehydration.
Examples of Heterogeneous Catalysis
1. Contact Process
2SO2(g) + O2(g)
2SO3(g)
The catalyst for the reaction is vanadium(V) oxide, V2O5. During the reaction, vanadium
changes its oxidation state; it is first reduced to vanadium(IV) by sulphur(IV) oxide and is
then oxidised back to vanadium(V) by oxygen.
SO2 + V2O5
SO3 + V2O4
2V2O4 + O2
2 V2O5
The oxidation state of vanadium is unchanged at the end of the reaction. This two-step
process has a lower activation energy than the uncatalysed route.
2. Haber Process
N2(g) + 3H2(g)
2NH3(g)
The catalyst in the Haber process is iron.
3. Manufacture of methanol
CO + 2H2
CH3OH
The catalyst is Cr2O3 and the reaction takes place at 300atm and 300oC.
TOPIC 13.23: THE TRANSITION METALS
20
Homogeneous Catalysis
Homogeneous catalysis usually takes place in solution. When transition metal ions act
as homogeneous catalysts, a change in oxidation state is usually involved, as the
reaction proceeds via the formation of an intermediate species.
An example of homogeneous catalysis is the oxidation of iodide ions by
peroxodisulphate(VI) ions. The reaction is catalysed by either iron(II) or iron(III) ions.
Despite a very favourable redox potential:
2I-(aq) + S2O82-(aq)
I2(aq) + 2SO42-(aq)
Eo = +1.51V
the uncatalysed reaction is slow since it has to occur by the collision of two negativelycharged ions. In the presence of the catalyst, the activation energy is lower, because the
reaction occurs by the collision of a negative ion with a positive ion.
When the reaction is catalysed by an iron(II) catalyst, the iron(II) ions are first oxidised to
iron(III) by peroxodisulphate(VI) ions. The iron(III) ions are then reduced back to iron(II)
ions by iodide ions.
2Fe2+(aq) + S2O82-(aq)
2Fe3+(aq) + 2SO42-(aq)
2I-(aq) + 2Fe3+(aq)
I2(aq) + 2Fe2+(aq)
When the reaction is catalysed by an iron(III) catalyst, the two steps take place in
reverse order. The iron(III) ions are first reduced to iron(II) ions by iodide ions; the iron(II)
ions are then oxidised back to iron(III) by peroxodisulphate(VI) ions.
2I-(aq) + 2Fe3+(aq)
I2(aq) + 2Fe2+(aq)
2Fe2+(aq) + S2O82-(aq)
2Fe3+(aq) + 2SO42-(aq)
Autocatalysis
In some reactions, one of the products of the reaction acts as a catalyst for the reaction.
This is called autocatalysis. In such reactions, the reaction rate increases initially as the
product is formed, instead of following the normal pattern of decreasing steadily. A plot
of reaction rate against time shows a curve with a maximum. A plot of % reaction
against time gives a characteristic sigmoid (s-shaped) curve.
Reaction
rate
% reaction
Time
TOPIC 13.23: THE TRANSITION METALS
Time
21
Examples:
1. Oxidation of ethanedioate ions by manganate(VII) ions.
This reaction can be carried out as a titration at about 70 oC. It is catalysed by the
Mn2+(aq) ions, which are formed in the reaction. The overall reaction is:
2MnO4- + 16H+ + 5C2O42-
2Mn2+ + 10CO2 + 8H2O
When potassium manganate(VII) solution is first added from the burette to an acidified
solution containing ethanedioate ions, the purple colour of manganate(VII) is not
discharged immediately. However, once this colour has been discharged, forming
Mn2+(aq) ions, the reaction is catalysed and further addition of manganate(VII) leads to
an immediate discharge of the colour, which continues until the end-point is reached.
The uncatalysed reaction is slow since it has to occur by the collision of two negativelycharged ions. In the presence of the catalyst, the activation energy is lower, because the
reaction occurs, in two stages, by the collision of a negative ion with a positive ion.
MnO4- + 8H+ +4Mn2+
5Mn3+ + 4H2O
2Mn3+ + C2O42-
TOPIC 13.23: THE TRANSITION METALS
2Mn2+ + 2CO2
22