Transition Metals

A transition metal is a metal atom that has a partially filled d sub-orbital in an atom or at-least 1 of its
stable ions. The transition metals are as follows:
However, Zn is not a transition metal as neither the Zn atom or Zn2+ ion have a partially filled d sub-orbital.
Similarly, the Sc3+ ion does not have a partially filled d sub-orbital so is not considered a transition metal.
When writing the electron configurations for these transition metals, write out the electron configuration
for the atom first (remembering that the 4s orbital fills before the 3d except in Cr and Cu where 1 electron
is promoted into the 3d sub-orbital). Then remove the electrons (from the 4s orbital first) to find what is
General properties of transition metals;
All transition metals have a partially filled d sub-orbital. They form coloured complexes because of this.
Transition metals also have variable oxidation states and catalytic behaviour.
Transition metal ions can combine with ligands (molecules or ions which donate an electron lone pair) to
form complexes via co-ordinate covalent bonds to the central metal ion.
 Some ligands from just 1 co-ordinate covalent bond -> Unidentate ligands; Water, NH3, Cl-
, CN

Some ligands from 2 co-ordinate covalent bonds -> Bidentate ligands; Ethane-1,2-diamine (‘en’)
and Ethanedioate ion C2O4
 Some ligands from more than 2 co-ordinate covalent bonds -> Multidentate ligands; The example
you have to know is EDTA which is used at a high pH so all carboxylic acid protons dissociate
forming [EDTA]4-
Co-ordination number;
The number of co-ordinate bonds that are formed between a metal ion and its ligands is known as the coordination number;
Ligands that can form more than 2 co-ordinate bonds are known as chelating agents. These are often used
when an individual is suffering from heavy metal poisoning as they wrap themselves around the metal
cation. EDTA is a good example.
With small ligands (Water, Ammonia, Cyanide etc…) they will usually form an octahedral shape, but with
large ligands (Cl-
ions) they will usually form a tetrahedral shape as not as many ligands can fit around the
central transition metal cation.
Iron (Fe2/3+) and Copper (Cu2+) will form octahedral complexes unless a large ligand is used;
All Platinum (Pt2+) and most Ni2+ ions will form square planar complexes (e.g. cis-platin);
Silver (Ag+
) will always form linear complexes;
Colorimetry can be used to determine the ratio of ligands to transition metal ion. Different solutions of
transition metal ion solution and ligand solutions are placed in a colorimeter, the solution that has the
highest absorbance will be the correct stoichiometric ratio of transition metal ion : ligand solution. Any
excess of either of these solutions results in dilution of the solution and the absorbance will decrease.
Some molecules (such as Hb) will form complexes with Oxygen atoms allowing it to be carried around the
body. However, molecules of CO will form permanent dative covalent bonds which are much stronger than
those formed by Oxygen, so will prevent the amount of Oxygen binding sites that are available.
Formation of coloured ions;
Many transition metal complexes are coloured thanks to their partially filled d sub-orbitals;
 Cobalt (II) Hexaaqua -> [Co(H2O)6]
2+ Pink
 Cromium (III) Hexaaqua -> [Cr(H2O)6]
3+ Green
 Iron (II) Hexaaqua -> [Fe(H2O)6]
2+ Pale Green
 Iron (III) Hexaaqua -> [Fe(H2O)6]
3+ Yellow
 Copper (II) Hexaaqua -> [Cu(H2O)6]
2+ Blue
 Copper (II) Chloride -> [CuCl4]
2- Yellow
 Cobalt (II) Chloride -> [CoCl4]
2- Blue

Silver (I) Diammonium -> [Ag(NH3)2]
Colours are dependent upon the oxidation state of the metal ion. As the oxidation state changes, so will
the colour. These colours arise when visible light is absorbed by electrons which are promoted from the
ground state to a higher energy level. There are 5 d sub-orbitals which are repelled by the 6 (or 4 or 2)
ligands when they approach along the x, y and z axes. As 2 of these orbitals are directly along the axes and
3 are in between the axes, we get the following splitting pattern;
The splitting in the energy levels is vital to the colour that is produced by
the transition metal cation.
The difference between the ground state and the higher energy level (ΔE) can be linked by the equation;
Where h is plank’s constant and v is the frequency of the photon. The light transmitted through the
solution consists of the colours that are not absorbed by the electrons in the ground state. If a metal atom/
ion has a completely empty or full d sub-orbital it is not capable of having electron transmissions, and so it
will appear colourless.
Higher the concentration of transition metal ions in solution, the more vibrant the colour will be (higher
absorbance of visible light in UV/Visible spectroscopy).
Variable oxidation states;
Many metals can have variable oxidation states (Fe can be II or III, Cu can be I or II)
Vanadium exists in 4 different oxidation states. Ammonium Vanadate can be dissolved in HCl to form the
first species which can then be reduced using Zn:
Oxidation State Ion Colour
+5 VO2
+ Yellow You
+4 VO2+ Blue Better
+3 V
3+ Green Get
+2 V
2+ Violet Vanadium
Similarly, adding HCl to potassium chromate and then adding Zn will reduce Chromium through its
oxidation states
Oxidation State Ion Colour
+6 K2Cr2O4 Yellow You
+6 Cr2O7
2- Orange Old
+3 Cr3+ Green Greedy
+2 Cr2+ Blue Boy
These types of compounds can be
used in redox titrations because a
colour change occurs when they are
oxidised or reduced. In exam questions
they like to use dichromate and
manganate ions

Catalytic behaviour;
Catalysts are substances that increase the rate of a reaction, by providing an alternative chemical pathways
with a lower Ea, without being used up or changed in the process.
Homogeneous Catalysis;
These catalysts are in the same phase as the reactants. Common examples include acid catalysts and AlCl3
dissolved in ethanol. Typically, homogeneous catalysts form reactive intermediates during the reaction
which increase the rate of a reaction by making one of the reactants more reactive (usually through
protonation). With homogeneous transition metal catalysts, the variable oxidation states allows for the
formation of reactive intermediates very easily.
Thiosulfate and Iodide ions;
This reaction normally occurs very slowly as the 2 negative ions repel each other, so we use an Iron
(Fe2+/Fe3+) catalyst;
These 2 equations can occur either way around and can be deduced from the E standard values if given
This is where the catalyst is formed as a product of the reaction. The initial rate is very slow, but as the
catalyst is formed, the rate will rapidly increase. An example is Ethanedioate ions and Manganate ions;
The Mn2+ manganese ion is the catalyst for this reaction and will increase the rate of the reaction in the
following way;

Heterogeneous Catalysis;
Catalysts are in a different phase to the reactants. Usually, these catalysts are solid metals. Reacting
molecules are adsorbed onto the surface of the catalyst at the active site. The reaction then occurs on the
surface of the catalyst. In order to be an effective catalyst, reactants must be able to move around the
surface of the catalyst between active sites to react with other reactant molecules. Once the reaction has
occurred, the molecule desorbs from the surface of the catalyst.
Adsorption increases the rate by;
 Bringing reactant particles closer together in the gas phase
 Weakens the bonds in the molecule, decreasing the Ea
 Hold molecules in the correct orientation to react
However, adsorption must be just the right strength. Too weak (Ag) and not many molecules are adsorbed
and too strong (W) means that molecules cannot move between active sites, so less likely to meet other
reactants. Ni and Pt have just the correct adsorption strength to be a good catalyst.
Catalysts should have a honeycomb structure to increase the surface area so that less catalyst is actually
needed. The catalyst is usually spread very thinly onto a support medium of ceramic metals to maximise
the S.A. and reduce costs.
Certain atoms may adsorb too strongly onto the active sites of the catalyst, blocking the active sites, this is
called poisoning. This decreases the surface area available for the reaction and so lowers the efficiency of
the catalyst. Often, poisons are very difficult to remove and so the catalyst must be replaced which is very
Poisons include lead in petrol which poisons catalytic converters and Sulfur in the Haber process which
poisons the Fe catalyst.
Contact Process;
This is used to produce SO3 which is used in manufacturing H2SO4. Vanadium Oxide (V2O5) is used as a

+ 4Mn2++ 8H+
-> 5Mn3+ + 4H2O
2Mn3+ + 2C2O4
2- -> 2Mn2+ + 2CO2
V2O5 + SO2 -> V2O4 + SO3
V2O4 + ½ O2 -> V2O5

Aqueous Ions in Solution;
When ligands bond to transition metal ions, the ligands are acting as Lewis bases (electron pair donators)
and transition metal ions act as Lewis acids. Aqueous ions in solution can be hydrolysed, effectively losing a
As 3+ metal ions are smaller and more highly charged than 2+ ions, it has a higher electron charge density,
so can polarise the OH group to a greater extent than the 2+ ions. This results in a weaker bond in the OH
group and so 3+ ions are more acidic than 2+ ions. Therefore, the equilibrium of a 3+ ion as shown above
lies further to the right than the 2+ ion.
Transition metal ions with the 2+ oxidation state;
Fe2+ Co2+ Cu2+
NaOH Fe(OH)2 precipitate formed
by hydrolysis; DARK GREEN
No further reaction with
excess NaOH
Co(OH)2 precipitate formed by
hydrolysis; DARK BLUE
No further reaction with excess
Cu(OH)2 precipitate formed
by hydrolysis; LIGHT BLUE
No further reaction with
excess NaOH
NH3 Fe(OH)2 precipitate formed
by hydrolysis; DARK GREEN
No further reaction with
excess NH3
Co(OH)2 precipitate formed by
hydrolysis; DARK BLUE
Reaction with excess NH3 to
form a straw coloured solution
by ligand substitution
Fe(OH)2 precipitate formed
by hydrolysis; LIGHT BLUE
Reaction with excess NH3 to
form a DARK BLUE solution
by ligand substitution
2- FeCO3 precipitate formed
by hydrolysis; DARK GREEN
No further reaction with
excess CO3
CoCO3 precipitate formed by
hydrolysis; PINK
No further reaction with excess
CuCO3 precipitate formed by
hydrolysis; GREEN/BLUE
No further reaction with
excess CO3

Transition metal ions with the 3+ oxidation state;
Fe3+ Cr3+ Al3+
NaOH [Fe(OH)3(H2O)3] precipitate
formed by hydrolysis;
No further reaction with
excess NaOH
[Cr(OH)3(H2O)3] precipitate
formed by hydrolysis; GREEN
Reacts with excess NaOH to
form [Cr(OH)6]
3- which is a
[Al(OH)3(H2O)3] precipitate
formed by hydrolysis; WHITE
Reacts with excess NaOH to
form [Cr(OH)4(H2O)2]
-which is
NH3 Fe(OH)3 precipitate formed
by hydrolysis; BROWN
No further reaction with
excess NH3
Cr(OH)3 precipitate formed by
hydrolysis; GREEN
Reaction with excess NH3 to
ligand substitution
Al(OH)3 precipitate formed
by hydrolysis; WHITE
No further reaction with
excess NH3
2- [Fe(OH)3(H2O)3] precipitate
formed by hydrolysis;
[Cr(OH)3(H2O)3] precipitate
formed by hydrolysis; GREEN
[Al(OH)3(H2O)3] precipitate
formed by hydrolysis; WHITE
Because the 3+ metal ions are more acidic, they will form a hydroxide instead of a carbonate. Also CO2 will
be formed which can be identified by effervescence.
Ligand Substitution Reactions;
Occurs when one ligand in a complex is replaced by another. This usually occurs with ammonia and Cl-
when in excess;
With ammonia, all (or most) of the water ligands are displaced but the co-ordination number does not
However, when Cl-
ions substitute for water molecules, they are much larger and so only 4 Cl-
ions are
substituted onto the transition metal cation.

This can also occur with EDTA in what is known as a chelating effect. EDTA forms a highly stable complex
because the substitution results in a large increase in entropy (so ΔG is very negative giving a
thermodynamically stable substance) and so the reverse reaction is not thermodynamically feasible as
there is a large increase in the number of molecules.