What are transition metals?
Lecture 1. CHEM1902 Coordination Chemistry
The elements of the second and third rows of the
show gradual changes in properties across the table from left to
right as expected. Electrons in the outer shells of the atoms of
these elements have little shielding effects resulting in an
increase in effective nuclear charge due to the addition of
protons in the nucleus. Consequently, the effects on atomic
properties are: smaller atomic radius, increased first ionization
energy, enhanced electronegativity and more nonmetallic
character. This trend continues until one reaches calcium (Z=20).
There is an abrupt break at this point. The next ten elements
called the first transition series are remarkably similar in
their physical and chemical properties. This general similarity
in properties has been explained in terms of their relatively
small difference in effective nuclear charge over the series.
This occurs because each additional electron enters the
penultimate 3d shell providing an effective shield between the
nucleus and the outer 4s shell.
Thus, the transition elements can be defined as those in which
the d electron shells are being filled and so we generally ignore
Sc and Zn where Sc(III) is d0 and Zn(II) is
Summary of Physical Properties
It is useful, at the beginning, to identify the physical and
chemical properties of transition elements which differ from main
group elements (s-block) such as Calcium.
- have large charge/radius ratio;
- are hard and have high densities;
- have high melting and boiling points;
- form compounds which are often paramagnetic;
- show variable oxidation states;
- form coloured ions and compounds;
- form compounds with profound catalytic activity;
- form stable complexes.
The following table summarises some of the physical properties of
/ kJ mol-1
||-engines/aircraft industry-density is 60% of iron
||-stainless steel, 19% Cr, 9% Ni the rest Fe
||-alloys eg with C steel, the most significant use
||-alloys eg with Cu
||-alloys eg with C steel, the most significant use
||-alloys eg with Cr and W for hardened drill bits
||-alloys Fe/Ni armour plating, resists corrosion
||-high electrical conductivity (2nd to Ag), wiring
Densities and metallic radii
The transition elements are much denser than the s-block elements
and show a gradual increase in density from scandium to copper.
This trend in density can be explained by the small and irregular
decrease in metallic radii coupled with the relative increase in
Melting and boiling points
The melting points and the molar enthalpies of fusion of the
transition metals are both high in comparison to main group
elements. This arises from strong metallic bonding in transition
metals which occurs due to delocalization of electrons facilitated
by the availability of both d and s electrons.
In moving across the series of metals from scandium to zinc a small change
in the values of the first and second ionization energies is observed.
This is due to the build-up of electrons in
the immediately underlying d-sub-shells that efficiently shields
the 4s electrons from the nucleus and minimizing the increase
in effective nuclear charge from element to element. The
increases in third and fourth ionization energy values are more
rapid. However, the trends in these values show the usual
discontinuity half way along the series. The reason is that the
five d electrons are all unpaired, in singly occupied orbitals.
When the sixth and subsequent electrons enter, the electrons have
to share the already occupied orbitals resulting in
inter-electron repulsions, which would require less energy to
remove an electron. Hence, the third ionization energy curve for
the last five elements is identical in shape to the curve for the
first five elements, but displaced upwards by about 580 kJ mol-1.
The electronic configuration of the atoms of the first row
transition elements are basically the same. It can be seen in the
Table above that there is a gradual filling of the 3d orbitals
across the series starting from scandium. This filling is,
however, not regular, since at chromium and copper the population
of 3d orbitals increase by the acquisition of an electron from
the 4s shell. This illustrates an important generalization about
orbital energies of the first row transition series. At chromium,
both the 3d and 4s orbitals are occupied, but neither is
completely filled in preference to the other. This suggests that
the energies of the 3d and 4s orbitals are relatively close for
atoms in this row.
In the case of copper, the 3d level is full, but only one electron
occupies the 4s orbital. This suggests that in copper the 3d
orbital energy is lower than the 4s orbital. Thus the 3d orbital
energy has passed from higher to lower as we move across the
period from potassium to zinc. However, the whole question of
preference of an atom to adopt a particular electronic
configuration is not determined by orbital energy alone. In
chromium it can be shown that the 4s orbital energy is still
below the 3d which suggests a configuration [Ar]
3d44s2. However due to the effect of
electronic repulsion between the outer electrons the actual
configuration becomes [Ar]3d54s1 where all
the electrons in the outer orbitals are unpaired. It should be
remembered that the factors that determine electronic
configuration in this period are indeed delicately balanced.
This shows that elemental Mn is a stronger reductant than
dihydrogen and hence should be able to displace hydrogen gas from
1 mol dm-3 hydrochloric acid.
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The Department of Chemistry, University of the West Indies,
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Created January 2004. Links checked and/or last
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