Edexcel A2 Unit 5 Transition Metal Properties

1. Transition Metals & Their Chemistry
Transition metals are elements which form one or more
stable ions with an incomplete d-subshell. Copper and zinc
are not transition elements.
D block elements are those elements which have electrons
filling in their d orbitals after the 4s orbitals…. I think.
Properties:
 Transition metals show variable oxidation states.
This is because the 4s and 3d orbitals have a
similar energy level. After electrons have been lost
from the 4s orbitals, some or all electrons may be
lost from the 3d orbitals. Electrons can transition
within the orbitals e.g Mn can promote one
electron to the 4p orbital to gain an oxidation state
of +7. For any covalent bond to be formed with the
central transition metal ion there must be an
unpaired electron present. Electrons can never be
lost from the 3p orbitals because it is at a much
lower energy level than the 3d-4s orbitals. The
minimum oxidation state is +1 (Cu) and the
maximum oxidation state is +7 (Mn). As a general
rule, the number of oxidation states a metal can
have decrease after Fe because in Fe the electron
pairing starts in the 3d orbital. The space to
promote or accept electrons decreases.

2.  They are often paramagnetic. Those elements only
which have an unpaired electron available. The
movement of an unpaired electron produces a
small magnetic field. When this is exposed to an
external magnetic field, there is a positive
interaction called paramagnetism. The lines of
forces become more concentrated and there is a
considerable increase in the magnetic field.
Elements which don’thave an unpaired electron
can not produce this effect and are called
diamagnetic.
 They are solids at room temperature with an
exception of Mercury.
 They have a silvery-blue colour exceptCu and Au.
 They have high densities, high conductivity, high
mp and bp. This is because they can delocalise 3d
and 4s electrons increasing the number of
electrons in the sea than normal. The size of the
cation further decreases due to greater effective
nuclear charge. The forces of attraction are
stronger, the lattice is more tightly packed and
there are more atoms per unit volume. Copper is
the exception along with half-full orbital elements.
Copper has full 3d and 4s orbitals so doesn’tlet go
of its electrons easily. There are fewer delocalised
electrons, the atomic/ionic size is larger so there

3. are less atoms per unit volume. Resulting in a
lower mp, bp, density and conductivity.
 They form coloured compounds. Hydrated d-block
elements are coloured where they have an
unpaired d-orbital electron available. Elements
which have configuration containing 3d0 or 3d10
they don’t form coloured ions. When a ligand
approaches the cation, the d-orbitals are split into
2 levels. Three orbitals are at a lower energy level
and 2 are at a high-energy level. This is called dd-
transition. When light is shone on the mixture, a
particular frequency from the light is absorbed
effectively removing it. An electron jumps from a
lower energy level to a higher energy level. When
this excited ion collides with another ion, the
excited electron falls back to the lower energy
level. The ion absorbs light again and the process
continues. The complementary colour of the light
absorbed is seen. Different ligands have different
splitting powers. Some split the orbitals more with
a greater energy difference, the higher frequency
of light is absorbed then and the longer the
wavelength of light seen. The splitting powers of
different ligands in the increasing order are:
chloride<fluoride<
hydroxide<water<ammonia<cyanide. Zinc forms

4. colourless or white solutions because it has a full
3d orbital so electrons can not undergo dd-
transition.
 They act as catalysts. Catalysts provide an
alternate reaction pathway with a lower activation
energy. This increases the number of particles with
energy greater than or equal to activation energy
and the rate of reaction increases. The partially
filled 3d orbitals of the metal can accept electrons
from the reactant molecules and change their
oxidation state. Catalysts are poisoned when
molecules other than the reactant molecules bind
irreversibly or more strongly in the active sites of
the catalyst. This decreases the number of active
sites available so reaction rate slows down.
 They have different types of bonds with different
oxidation states. +1, +2 and +3 form ionic bonds.
+4 and above form covalent bonds producing
anions. Anhydrous chlorides, bromides, iodides
form covalent bonds with the central metal ion
(hydrated sometimes become ionic).
 Different cation charges are possible because
hydration or lattice enthalpy can compensatefor
the energy required to remove and extra electron
e.g. Fe+2 to Fe+3. Group 1 and 2 elements can not
provide for this extra energy.

5.  The atomic radius and the ionic radius slightly
decrease across the period. Electrons are being
added to inner 3d orbitals so electron shielding
doesn’tincrease much. A greater effective nuclear
charge is experienced by the increase in protons.
 The first ionisation energies are almost similar
because the electrons are being removed from the
4s orbitals. Although they are different. The second
and the third ionisation energies increase
respectively across the period because of the
decrease in electron shielding and the increase in
experience of effective greater nuclear charge.
 The E-Naught values for M+2 to M are negative
except for copper so they react with hydrochloric
acid to produce hydrogen.
 They form interstitial compounds with carbon,
nitrogen, boron and hydrogen because their
atomic radius is small enough to fit in the lattice of
the metal.
 These elements can form alloys with each other
because they have similar atomic size so can fit in
each other’s lattices without disrupting it much.
 They can form complex ions. When d-block cations
are dissolved in water, they become hydrated. This
is because the lone pair of electrons on the oxygen
atom of the water bonds with an empty d-orbital

6. of the cation. A similar process happens with all
ligands.
Complex Ions & Their Types:
 The molecules forming a dative covalent bond with
the central metal ion are called ligands.
 The co-ordination number is the number of ligands
attached to metal ion.
 The complex ion is the whole thing.
 Each ligand acts as a Lewis base by donating a pair
of electrons and as bonds are being formed, energy
is evolved.
 Monodentateligands:Each ligand bonds using one
lone pair and forms one dative covalent bond.
 Bidentateligands: Each ligand has 2 lone pairs of
electrons and forms 2 dative bonds per ligand.
 Polydentateligands: They have more than 2 lone
pairs on each ligand.
 Neutral molecules e.g. water and ammonia can fit
in an octahedral shape around the central ion i.e.
the coordination number is 6. The bond angle is 90
degrees.

7.  Negative ligands e.g. chloride or cyanide can’t fit
more than 4 ligands around a metal ion due to
electron repulsion and also the negative ion is
larger than the metal ion. The shape is tetrahedral
and the bond angle is 109.5 degrees. The
coordination number is 4.
 Ligands like EDTA-4
(has 6 lone pairs) cannot fit
more than one ligand around the ion due to its
size.
 When the coordination number is 2 then the
complex has a linear shape e.g. Diamminesilver(I)
ion which is Tollens’s reagent.
 When the central ion has configuration of d8
orbital then a square planar complex is formed
with 4 ligands. These ligands show
stereoisomerism. They have cis-trans isomers. e.g.
(Pt(NH3)(Cl)2). Cis-platin is an anti-cancer drug.
They chloride ions can easily be lost and replaced
and the complex bonds with the nitrogen atom of
the DNA in the cancerous cell. This prevents the
cell from reproducing and the cell dies.

8.  Stereoisomerismarises when a bidentateligand
uses all of its 6 lone pairs to bond with the ion.
Non-superimposable mirror images are formed.