Download THE GENERAL FEATURES OF TRANSITION METAL CHEMISTRY

THE GENERAL FEATURES OF
TRANSITION METAL
CHEMISTRY
General features of transition metals
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This presentation explains what a transition
metal is in terms of its electronic structure,
and then goes on to look at the general
features of transition metal chemistry.
These include
variable oxidation state (oxidation number)
complex ion formation
coloured ions
magnetic properties
catalytic activity
What is a transition metal?
• The terms transition metal (or element)
and d block element are sometimes
used as if they mean the same thing.
• They don't - there's a subtle (not
obvious) difference between the two
terms.
d block elements
• You will remember that when you are
building the Periodic Table and working
out where to put the electrons,
something odd happens after argon.
• At argon, the 3s and 3p levels are full,
but rather than fill up the 3d levels next,
the 4s level fills instead to give
potassium and then calcium.
• Only after that do the 3d levels fill.
d block elements
The elements in the Periodic Table which
correspond to the d sublevels filling are called d
block elements. The first row of these is shown in
the shortened form of the Periodic Table below.
The electronic structures of the d block elements shown are:
Sc
[Ar] 4s2 3d1
III B
Ti
[Ar] 4s2 3d2
IVB
V
[Ar] 4s2 3d3
VB
Cr
[Ar] 4s1 3d5
VIB
Mn
[Ar] 4s2 3d5
VIIB
Fe
[Ar] 4s2 3d6
VIIIB
Co
[Ar] 4s2 3d7
VIIIB
Ni
[Ar] 4s2 3d8
VIIIB
Cu
[Ar] 4s1 3d10
IB
Zn
[Ar] 4s2 3d10
IIB
You will notice that the pattern of filling isn't entirely tidy! It is
broken at both chromium and copper.
Cr
[Ar] 3d54s1 instead of [Ar] 3d44s2
 This is something that you are just going to have to accept.
 There is no simple explanation for it which is usable at this
level.
 Any simple explanation which is given is faulty!
 People sometimes say that a half-filled d level as in
chromium (with one electron in each orbital) is stable, and
so it is - sometimes!
 But you then have to look at why it is stable.
 The obvious explanation is that chromium takes up this
structure because separating the electrons minimizes the
repulsions between them - otherwise it would take up some
quite different structure.
 But you only have to look at the electronic configuration of
tungsten (W) to see that this apparently simple explanation doesn't
always work.
 Tungsten has the same number of outer electrons as chromium, but
its outer structure is different - 5d46s2. Again the electron repulsions
must be minimized - otherwise it wouldn't take up this
configuration.
 But in this case, it isn't true that the half-filled state is the most
stable - it doesn't seem very reasonable, but it's a fact!
 The real explanation is going to be much more difficult than it
seems at first sight.
 Neither can you use the statement that a full d level (for example, in
the copper case) is stable, unless you can come up with a proper
explanation of why that is. You can't assume that looking nice and
tidy is a good enough reason!
 If you can't explain something properly, it is much better just to
accept it than to make up faulty explanations which sound OK on
the surface but don't stand up to scrutiny!
Transition metals
• Not all d block elements count as
transition metals! There are
discrepancies between the various
syllabuses, but the majority use the
definition:
• A transition metal is one which forms
one or more stable ions which have
incompletely filled d orbitals.
SCANDIUM , ZINC and COPPER
 On the basis of this definition, scandium and zinc don't count as
transition metals - even though they are members of the d block.
 Scandium has the electronic structure [Ar] 3d14s2. When it forms
ions, it always loses the 3 outer electrons and ends up with an argon
structure. The Sc3+ ion has no d electrons and so doesn't meet the
definition.
 Zinc has the electronic structure [Ar] 3d104s2. When it forms ions, it
always loses the two 4s electrons to give a 2+ ion with the electronic
structure [Ar] 3d10. The zinc ion has full d orbitals and doesn't meet
the definition either.
 By contrast, copper, [Ar] 3d104s1, forms two ions. In the Cu+ ion the
electronic structure is [Ar] 3d10. However, the more common Cu2+ ion
has the structure [Ar] 3d9.
 Copper is definitely a transition metal because the Cu2+ ion has an
incomplete d orbitals.
Transition metal ions
 You have already come across the fact that when the
Periodic Table is being built, the 4s orbital is filled
before the 3d orbitals.
 This is because before filling orbitals, 4s orbitals
have a lower energy than 3d orbitals.
 However, once the electrons are actually in their
orbitals, the energy order changes - and in all the
chemistry of the transition elements, the 4s orbital
behaves as the outermost, highest energy orbital.
Transition metal ions
 The reversed order of the 3d and 4s orbitals only
applies to building the atom up in the first place. In all
other respects, you treat the 4s electrons as being
the outer electrons.
 This is another of those things that you just have to
accept.
 The explanation again lies well beyond the level you are
working at. Just remember that once you have the full
electronic structure for one of these atoms, the 4s
electrons are the outermost electrons.
Remember this:
 When d-block elements form ions, the 4s electrons
are lost first.
Transition metal ions
To write the electronic structure for Co2+:
Co
[Ar] 3d74s2
Co2+
[Ar] 3d7
The 2+ ion is formed by the loss of the two 4s electrons.
To write the electronic structure for V3+:
V
[Ar] 3d34s2
V3+
[Ar] 3d2
The 4s electrons are lost first followed by one of the 3d
electrons.
Transition metal ions
To write the electronic structure for Cr3+:
Cr
Cr3+
1s22s22p63s23p63d54s1
1s22s22p63s23p63d3
The 4s electron is lost first followed by two of the 3d
electrons.
To write the electronic structure for Zn2+:
Zn
Zn2+
1s22s22p63s23p63d104s2
1s22s22p63s23p63d10
This time there is no need to use any of the 3d electrons.
To write the electronic structure for Fe3+:
Fe
Fe3+
1s22s22p63s23p63d64s2
1s22s22p63s23p63d5
Forming transition metal ions
The rule is quite simple:
Take the 4s electrons off first, and then
as many 3d electrons as necessary to
produce the correct positive charge.
Variable oxidation state (number)
 One of the key features of transition metal chemistry is
the wide range of oxidation states (oxidation numbers)
that the metals can show.
 It would be wrong, though, to give the impression that
only transition metals can have variable oxidation
states.
 For example, elements like sulphur or nitrogen or
chlorine have a very wide range of oxidation states in
their compounds - and these obviously aren't transition
metals.
 However, this variability is less common in metals
apart from the transition elements.
 Of the familiar metals from the main groups of the
Periodic Table, only lead and tin show variable
oxidation state to any extent.
Examples of variable oxidation states in the transition
metals
Iron
Iron has two common oxidation states (+2 and +3) in, for
example, Fe2+ and Fe3+. It also has a less common +6
oxidation state in the ferrate(VI) ion, FeO42-.
Manganese
Manganese has a very wide range of oxidation states in its
compounds. For example:
+2 in Mn2+
+3 in Mn2O3
+4 in MnO2
+6 in MnO42+7 in MnO4-
Explaining the variable oxidation states in
the transition metals
 We'll look at the formation of simple ions like Fe2+ and Fe3+.
 When a metal forms an ionic compound, the formula of the
compound produced depends on the energetics of the process. On
the whole, the compound formed is the one in which most energy is
released.
 The more energy released, the more stable the compound.
 There are several energy terms to think about, but the key ones are:
 The amount of energy needed to ionize the metal (the sum of the
various ionization energies)
 The amount of energy released when the compound forms.
 This will either be lattice enthalpy if you are thinking about solids, or
the hydration enthalpies of the ions if you are thinking about
solutions.
 The more highly charged the ion, the more electrons you have to
remove and the more ionization energy you will have to provide.
 But off-setting this, the more highly charged the ion, the more energy
is released either as lattice enthalpy or the hydration enthalpy of the
metal ion.
Thinking about a typical non-transition
metal (calcium)
Calcium chloride is CaCl2. Why is that?
 If you tried to make CaCl, (containing a Ca+ ion), the
overall process is slightly exothermic.
 By making a Ca2+ ion instead, you have to supply more
ionisation energy, but you get out lots more lattice
energy.
 There is much more attraction between chloride ions
and Ca2+ ions than there is if you only have a 1+ ion.
 The overall process is very exothermic.
 Because the formation of CaCl2 releases much more
energy than making CaCl, then CaCl2 is more stable and so forms instead.
Thinking about a typical non transition
metal (calcium)
 What about CaCl3? This time you have to remove yet
another electron from calcium.
 The first two come from the 4s level. The third one
comes from the 3p. That is much closer to the nucleus
and therefore much more difficult to remove. There is a
large jump in ionisation energy between the second and
third electron removed.
 Although there will be a gain in lattice enthalpy, it isn't
anything like enough to compensate for the extra
ionisation energy, and the overall process is very
endothermic.
 It definitely isn't energetically sensible to make CaCl3!
Thinking about a typical transition metal (iron)
Here are the changes in the electronic structure of iron to
make the 2+ or the 3+ ion.
Fe
[Ar] 3d64s2
Fe2+
[Ar] 3d6
Fe3+
[Ar] 3d5
The 4s orbital and the 3d orbitals have very similar energies.
There isn't a huge jump in the amount of energy you need to
remove the third electron compared with the first and second.
Thinking about a typical transition
metal (iron)
 The figures for the first three ionisation energies (in kJ mol-1)
for iron compared with those of calcium are:
metal 1st IE 2nd IE 3rd IE
Ca
590
1150
4940
Fe
762
1560
2960
 There is an increase in ionisation energy as you take more
electrons off an atom because you have the same number of
protons attracting fewer electrons.
Thinking about a typical transition metal (iron)
 However, there is much less increase when you take
the third electron from iron than from calcium.
 In the iron case, the extra ionisation energy is
compensated more or less by the extra lattice
enthalpy or hydration enthalpy evolved when the 3+
compound is made.
 The net effect of all this is that the overall enthalpy
change isn't vastly different whether you make, say,
FeCl2 or FeCl3. That means that it isn't too difficult to
convert between the two compounds.
Variable oxidation states
 The multiple oxidation states of the d-block (transition
metal) elements are due to the proximity of the 4s and
3d sub shells (in terms of energy).
 All transition metals exhibit a 2+ oxidation state (both
electrons being lost from the 4s and all have other
oxidation states (common).
Sc Ti V Cr Mn Fe Co Ni Cu Zn
+1
+2 +2 +2 +2 +2 +2 +2 +2 +2
+3 +3 +3 +3 +3 +3 +3 +3
+4 +4
+4
+5
+6 +6 +6
+7
Some examples of complex ions formed by
transition metals
[Fe(H2O)6]2+
[Co(NH3)6]2+
[Cr(OH)6]3[CuCl4]2 Other metals also form complex ions - it isn't
something that only transition metals do.
 Transition metals do, however, form a very wide
range of complex ions.
The formation of complex ions
What is a complex ion?
 A complex ion has a metal ion at its centre with a
number of other molecules or ions surrounding it.
These can be considered to be attached to the
central ion by co-ordinate (dative covalent) bonds. (In
some cases, the bonding is actually more
complicated than that.)
 The molecules or ions surrounding the central metal
ion are called ligands.
 Simple ligands include water, ammonia and chloride
ions.
 What all these have got in common is active lone
pairs of electrons in the outer energy level.
 These are used to form co-ordinate bonds with the
metal ion.
Bonding in simple complex ions
Al(H2O)6 3+
 We are going to look in detail at the bonding in the
complex ion formed when water molecules attach
themselves to an aluminium ion to give Al(H2O)63+.
 Start by thinking about the structure of a naked
aluminium ion before the water molecules bond to it.
 Aluminium has the electronic structure
1s22s22p63s23px1
 When it forms an Al3+ ion it loses the 3-level
electrons to leave
1s22s22p6
Bonding in simple complex ions
 That means that all the 3-level orbitals are now
empty. The aluminium uses of six of these to accept
lone pairs from six water molecules.
 It re-organises (hybridises) the 3s, the three 3p, and
two of the 3d orbitals to produce six new orbitals all
with the same energy.
 You might wonder why it chooses to use six orbitals
rather than four or eight or whatever.
 Six is the maximum number of water molecules it is
possible to fit around an aluminium ion (and most
other metal ions).
 By making the maximum number of bonds, it
releases most energy and so becomes most
energetically stable.
Only one lone pair is shown on each water molecule.
The other lone pair is pointing away from the
aluminium and so isn't involved in the bonding. The
resulting ion looks like this:
Dotted arrows represent lone pairs
coming from water molecules behind
the plane of the screen or paper.
Wedge shaped arrows represent
bonds from water molecules in front of
the plane of the screen or paper.
Al(H2O)6
3+
 Because of the movement of electrons towards the
centre of the ion, the 3+ charge is no longer located
entirely on the aluminum, but is now spread over the
whole of the ion.
 Because the aluminum is forming 6 bonds, the coordination number of the aluminum is said to be 6.
 The co-ordination number of a complex ion counts
the number of co-ordinate bonds being formed by
the metal ion at its centre.
 In a simple case like this, that obviously also counts
the number of ligands - but that isn't necessarily so,
as you will see later.
 Some ligands can form more than one co-ordinate
bond with the metal ion.
Fe(H2O)6
3+
This example is chosen because it is very similar to
the last one - except that it involves a transition metal.
Iron has the electronic structure
1s22s22p63s23p63d64s2
When it forms an Fe3+ ion it loses the 4s electrons and
one of the 3d electrons to leave
1s22s22p63s23p63d5
Looking at this as electrons-in-boxes, at the bonding
level:
Fe(H2O)6
+3
Looking at this as electrons-in-boxes, at the bonding
level:
 Now, be careful! The single electrons in the 3d level
are NOT involved in the bonding in any way.
 Instead, the ion uses 6 orbitals from the 4s, 4p and
4d levels to accept lone pairs from the water
molecules.
 Before they are used, the orbitals are re-organised
(hybridised) to produce 6 orbitals of equal energy.
Once the co-ordinate bonds have been formed, the ion
looks exactly the same as the equivalent aluminium ion.
Because the iron is forming 6 bonds, the co-ordination
number of the iron is 6.
CuCl4
2-
 This is a simple example of the formation of a
complex ion with a negative charge.
 Copper has the electronic structure
1s22s22p63s23p63d104s1
 When it forms a Cu2+ ion it loses the 4s electron and
one of the 3d electrons to leave
1s22s22p63s23p63d9
 To bond the four chloride ions as ligands, the empty
4s and 4p orbitals are used (in a hybridized form) to
accept a lone pair of electrons from each chloride ion.
 Because chloride ions are bigger than water
molecules, you can't fit 6 of them around the central
ion - that's why you only use 4.
Only one of the 4 lone pairs on each chloride ion is
shown. The other three are pointing away from the
copper ion, and aren't involved in the bonding.
CuCl42That gives you the complex ion:
 The ion carries 2 negative charges overall. That
comes from a combination of the 2 positive charges
on the copper ion and the 4 negative charges from
the 4 chloride ions.
 In this case, the co-ordination number of the copper
is, of course, 4.
Coordinated ligands
Ligands are the molecules (or ions) which donate an electron
pair to form a dative covalent bond with the central transition
metal atom (forming a complex molecule or ion).
Complexes
These are species which are formed around a central atom, with
other atoms, ions or molecules donating an electron pair to form
a covalent bond to this central atom. The result is a "complex"
usually an ion but may also be a molecule.
Complex
shape
ligands
coordination
name
number
3+
[Fe(H2O)6] octahedral water
6
hexa-aqua iron III ion
hexacyano ferrate III
[Fe(CN)6]3- octahedral cyanide CN6
ion
tetrachloro cuprate I
[CuCl4]3tetrahedral chloride Cl4
ion
square
tetra-ammine copper II
[Cu(NH3)4]2+
ammonia
4
planar
ion
+
[Ag(NH3)2] linear
ammonia
2
diammine silver I ion
carbon
tetracarbonyl Nickel 0
Ni(CO)4
tetrahedral
4
monoxide
molecule
Coloured compounds
 The color in the transition metals (d-block) is predominantly
due to the splitting of the d shell orbitals into slightly
different energy levels.
 As a result, certain wavelengths of energy can be absorbed
by the d-block elements (with electrons jumping between
these slightly different energy levels), resulting in the
complement color being visible.
 Colour is affected by both the oxidation state of the
transition metal and the type of ligand
Complex
Oxidation state of
ion
metal
[Fe(H2O)6]3+
III
[Fe(H2O)6]2+
II
[Cu(H2O)6]2+
II
[Cu(NH3)4]2+
II
[CuCl4]2-
II
colour
ligand
pale green
yellow
blue
deep blue
water
water
water
ammonia
chloride
ion
green
Non-transition metal ions
Transition metal ions
Magnetism
Transition metals and their ions often have unpaired 'd'
electrons which produce an asymmetric magnetic field
that can be detected. This is called paramagnetism
Examples
Complex
electronic
ion
configuration
[Fe(H2O)6]3+
[Ar]4s0 3d5
[Cr(H2O)6]3+
[Ar]4s0 3d3
[Cu(H2O)6]2+
[Ar]4s0 3d9
[Ni(NH3)6]2+
[Ar]4s0 3d8
[CoCl4]2[Ar]4s0 3d7
no of unpaired
electrons
5
3
1
2
3
magnetism
paramagnetic
paramagnetic
paramagnetic
paramagnetic
paramagnetic
Catalytic activity
 'd' block elements make good catalysts due to their
multiple oxidation states (hence their ability to react
with different species and produce a path of lower
activation energy, and so allow the reaction to proceed
at a faster rate).
 Another possible reason for their catalytic activity is
their available 'd' orbitals which allow reacting
molecules to co-ordinate to the surface of the transition
metal which in turn weakens the bonding within the
molecule allowing it to react.
Examples...




MnO2 in decomposition of hydrogen peroxide
V2O5 in the contact process
Fe in Haber process
Ni in conversion of alkenes to alkanes
REFERENCES
• http://www.chemguide.co.uk/inorganic/
complexions/colour.html
• http://ibchem.com/IB/ibnotes/brief/perhl.htm#oxid#oxid
• http://chemed.chem.wisc.edu/chempath
s/GenChem-Textbook/Transition-MetalIons-in-Aqueous-Solutions-1055.html
• http://www.wou.edu/las/physci/ch462/t
mcolors.htm