(viii) Transition Metal Configuration and Oxidation States

Transition Metals

A d-block transition metal has an incomplete d-subshell in at least one of their ions.

The filling of the d-orbitals follows the Aufbau principle - filling orbitals of lowest energy first. (4s before 3d)

Why do chromium [Ar] 3d⁵ 4s¹ and copper [Ar] 3d¹⁰ 4s¹ appear not to follow this rule?

Answer

However, whenever transition metals form ions, electrons are lost first from the outermost subshell, the 4s.

Oxidation States and Numbers

The oxidation state tells us the actual charge on an atom if it existed as a monatomic ion. The following rules are used to determine the oxidation state.

1. Uncombined elements have an oxidation number of 0.
2. The sum of all oxidation numbers in a molecule or neutral compound must equal zero.
3. The oxidation number of the atom in a polyatomic ion must add up to the charge of the ion.
4. Oxygen is -2, (except in peroxides where it is -1).
5. Hydrogen is +1, (except in metallic hydrides where it is -1 e.g. Li⁺H^-)
6. Group 7 halogens are -1, (except in compounds with oxygen or bonded to a halogen above itself in group 7 when they have a positive value).
7. Group 1 metals are always +1
8. Group 2 metals are always +2
9. Aluminium is always +3

Multiple Oxidation states

Many stable compounds exist with transition metals in different oxidation states. Compounds of the same transition metal in different oxidation states may have different colours.

The stability of transition metal oxidation states is determined by:

• electronic structure
• type of bonding
• stereochemistry

Oxidation can be defined as an increase in oxidation number. Reduction can be considered as a decrease in oxidation number.

The table below shows some common oxidation numbers of transition metals with 3d electrons:

Oxidation State	Metal Symbol
	Sc	Ti	V	Cr	Mn	Fe	Co	Ni	Cu	Zn
+1									x
+2		x	x	x	x	x	x	x	x	x
+3	x	x	x	x	x	x	x	x
+4		x	x		x
+5			x
+6				x	x	x
+7					x

Compounds containing metals in high oxidation states are often oxidising agents (e.g. permanganate, MnO₄⁻, where Mn is +7). Compounds containing metals in low oxidation states are often reducing agents (e.g. Fe²⁺).

Changes in oxidation number show whether oxidation or reduction has happened. e.g. MnO₄⁻ → Mn²⁺ is a decrease from +7 to +2, so manganese has been reduced.

(ix) Transition Metal Complexes

Ligands and Complexes

A complex consists of a central metal ion surrounded by ligands.

Ligands are electron donors and may be negative ions or neutral molecules but must contain an atom with lone pair (non-bonding pair) of electrons which can form a dative (co-ordinative) bond with a metal ion. (Remember in dative bonds, both electrons making up the covalent bond come from the same atom.)

If the ligand can only form one dative (co-ordinative) bond then it is classified as monodentate ('one toothed'). Common ones include the following;

Complete the table drawing the structure of each ligand showing the lone pair and charge where appropriate.

Ion	Name	Structure
chloride, Cl-	chlorido
water, H2O	aqua
ammonia, NH3	ammine
cyanide, CN-	cyanido

If the ligand can simultaneously form more than one dative bond it is called a polydentate ligand. Some important bidentate ('two toothed') ligands include,

Chelating ligands (chain-like) are polydentate ligands which are able to form more than one bond simultaneously with the same metal ion.

EDTA^4- - the ethylenediaminetetraacetate ion is a hexadentate ligand which forms very stable complexes with metals.

The number of bonds the central metal ion can form with ligands is known as the coordination number. For example, in [Cu(H₂O)₆]²⁺ the copper(II) is surrounded by six water molecules so will have a coordination number of 6. In [CuCl₄]^2- there are four chloride ions acting as ligands around the copper (II) ion and so the copper(II) ion in this complex has a co-ordination number of 4.

Naming Complexes

Complexes are written and named according to IUPAC rules.

1. Ligands are named in alphabetical order followed by the name of the metal and its oxidation state. If there is more than one of a ligand it is preceded by the prefix for the number di, tri, tetra etc.
2. If the ligand is a negative ion ending in -ide, for example chloride you drop the 'e' and add 'o' to become chlorido, cyanide becomes cyanido. Some common ligand names are given below. Note the names for ammonia and water.

3. If the complex ion is negative then the metal ion will have the suffix 'ate'. Sometimes latin names are used.

4. If the complex is a salt, the name of the positive ion (cation) precedes the name of the negative ion (anion).
5. The symbol of the metal is written first, followed by the symbols of the ligands in alphabetical order, with the atom that binds written first. For example water would be written OH₂ rather than H₂O. Oxalate would be, O₂C₂O₂ is used in the formula, rather than C₂O₄.
6. The formula of a complex should be enclosed in square brackets, with the charge outside the bracket e.g. [Fe(OH)₂(H₂O)₄]⁺ (Although you may see common complexes written without, like MnO₄^-)

Ligand names you need: water (aqua), ammonia (ammine), fluoride (fluorido), chloride (chlorido), bromide (bromido), iodide (iodido), cyanide (cyanido), hydroxide (hydroxido) and oxalate (oxalato).

In a negative complex ion, copper becomes cuprate and iron becomes ferrate, e.g. [CuCl₄]²⁻ is tetrachloridocuprate(II).

Transition Metals Test

Colour in Transition Metal Complexes - The Spectrochemical Series

Most transition metals form coloured complexes. The compounds are coloured because they absorb radiation from the visible part of the spectrum.

Complexes of transition metals can absorb light because the d-orbitals of the transition metal split and are no longer degenerate (of the same energy). When the metal is surrounded by ligands an electrostatic repulsion will raise the energy of any electron in orbitals pointing along these axes. Electrons in any orbitals pointing between axes will be less affected. Resulting in the previously degenerate orbitals (of the same energy) being split as shown.

The difference in energy between the two sets of orbitals (d-d transition) is called the crystal field strength (Δ).

Transition metal complexes are therefore coloured because they are able to absorb light when an electron is excited from the d-orbitals of the lower energy (ground state) to d orbitals of higher energy (excited state).

The energy difference between the sets of d orbitals depends on the position of the ligand in the spectrochemical series.

A 'stronger' ligand will displace a 'weaker' one from a metal ion.

Ligands that cause a large difference in energy between subsets of d orbitals are strong field ligands. Weak field ligands cause a small energy difference.

Colours of many transition metal complexes can be explained in terms of d-d transitions. Light is absorbed when electrons in a lower energy d orbital are promoted to a d orbital of higher energy.

If light of one colour is absorbed, the complementary colour is observed (see the colour wheel in (a)).

(x) UV and Catalysis

UV and Visible Spectroscopy

The effects of d-d transitions can be studied using ultra-violet and visible absorption spectroscopy. The wavelength for ultra-violet light is approximately 200-400 nm and 400-700 nm for visible.

Absorption of energy in the ultraviolet and visible regions of the spectrum causes transitions between electronic energy levels in atoms, molecules and ions. This energy difference corresponds to region of the electromagnetic spectrum involved and is greatest when ultraviolet is absorbed because ultraviolet radiation has a shorter wavelength (200 - 400 nm) and therefore higher energy than visible radiation.

Transition metal complexes are coloured because they absorb in the visible spectrum. If a compound absorbs solely in the ultraviolet it will appear colourless as our eyes are unable to detect that ultra violet radiation has been absorbed.

Ultra-violet/visible spectroscopy is a useful analytical tool. The sample used in solution is placed in a 'cell' cuvette. This is compared to an identical cell containing pure solvent. Radiation is scanned continuously through both samples and the spectrometer compares the two beams.

The absorption band for the cyan (blue-green) [Cu (H₂O)₆]²⁺ ion, found in aqueous solutions of copper(II) sulfate and copper(II) chloride is shown below.

The [Cu(H₂O)₆]²⁺ ion has an absorption band in the red part of the visible spectrum. Since red light is absorbed, blue and green light are transmitted which is why the aqueous copper(II) salts appear blue - green (cyan) to our eyes.

Catalysis

Transition metals or their compounds act as catalysts in many important biological and industrial chemical reactions.

Process	Catalyst Used
Haber Process	iron
Contact (to make sulfuric acid)	vanadium(V) oxide
Ostwald (to make nitric acid)	Platinum
Catalytic converter in car exhaust	platinum, palladium and rhodium
Preparation of methanol	copper
Manufacture of margarine	nickel
Polymerisation of alkene	titanium compounds

Many transition metals act as catalysts because of their ability to exist in a variety of different oxidation states.

The transition metal then reverts back (unchanged) to its original oxidation state once the reaction is complete.

Heterogeneous catalysts are in a different physical state to the reactants being catalysed. Heterogeneous catalysis can be explained in terms of the formation of activated complexes and the adsorption of reactive molecules onto active sites. The presence of unpaired d electrons or unfilled d orbitals is thought to allow activated complexes to form. This can provide reaction pathways with lower activation energies compared to the uncatalysed reaction.

Homogeneous catalysts are in the same physical state to the reactants being catalysed. Homogeneous catalysis can be explained in terms of changing oxidation states with the formation of intermediate complexes.

Transition Metals Colour and Catalysts Test