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Addition of ammonia to a solution containing [Cu(H₂O)₆]2+ causes 4 ammonia molecules to replace water molecules, forming a royal blue solution containing [Cu(NH₃)₄(H₂O)₂]²⁺ ions. This complex is octahedral, but as it contains 2 different ligands, there could be 2 different arrangements of ligands.

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These re tetrahedral complexes, all 4 chlorides at 109.5° to each other. The complexes are formed when copper(II) or cobalt(II) ions react with concentrated hydrochloric acid, which displaces the water molecules. There are distinct colour changes as the change in ligands & coordination geometry both contribute to changes in the light absorbed. Copper(II) goes from pale blue to yellow/green & cobalt(II) goes from pink to blue.

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These complexes are present in most aqueous solutions of Cu²⁺ & Co2+ giving the familiar colours of these solutions & many compounds.The complexes are octahedral with 1 lone pair on each oxygen atom of the water molecules used for bonding to the metal ion

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When a transition metal ions is exposed to a mixtrure of ligands, such as aq solutions containing chloride ions, ligands can be exchanged to form a new complex. This is an equilibrium process, so the concentrations of the metal ions & any possible ligands add may go identifying the species that will be present in solution.
[Cu(H₂O)₆]²⁺ + 4NH₃ ⇌ [Cu(H₂O)₂(NH₃)₄]²⁺ + 4H₂O
According to Le Chatelier’s principle, addition of extra ammonia shift the equilibrium to the left, producing more of [Cu(H₂O)₂(NH₃)₄]²⁺. Addition of extra water shifts the equilibrium to the right, producing more of the [Cu(H₂O)₆]²⁺ complex. This is associated with s colour change, with the ammonia-containing complex being royal blue compared to the pale blue of the orginal complex.
The equilibrium between complexes can lead to a change in geometry depending on the ligands used. The equilibrium below shows the interconversion between 2 complexes of cobalt. If a large amount of chloride is used, such as adding conc. HCl, then the equilibrium shifts from the octahedral complex to a tetrahedral chloro-complex.
[Co(H₂O)₆]²⁺ + 4Cl- ⇌ [CoCl₄]²⁻ + 6H₂O

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Most of the ions we have written as simple ions are actually complexes with water molecules as ligands around the transition metal atom or ion.

Both shapes can be seen for the same transition metal ion with different ligands. The shape found is dependent on the metal, the oxidation state of the metal and the ligands, and these factors often favour the octahedral complex with 6 ligands around the metal atom.

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Reactions of transition metal ions with hydroxide ions

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Reactions of transition metal ions with hydroxide ions

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Examples:

Iron (Fe) - The Haber process - to make ammonia (NH₃) from Nitrogen (N) and Hydrogen(H)

Nickel (Ni) - The hydrogenation of vegetable oils to for margarine

Platinum (Pt) - The oxidation of ammonia (NH₃) to form Nitric Acid (HNO₃)

Vanadium Oxide (V₂O₅) - The contact process for the production of sulfuric acid (H₂SO₄)

Mangansese Dioxide (MnO₂) - The catalytic decomposition of Hydrogen Peroxide (H₂O₂)

Transition metal catalysts are used in many industrial processes that would be uneconomical without them and are therefore essential. They are neede to make the majority of plastics, fertilisers explosives and some acids and solvents such as ethanoic acid and ethanal.

Heterogeneous Catalysts
In a different phase to the reactants (usually a solid catalyst and a gaseous reactant)

1. Reactants adsorb (bonded) onto surface (onto active sites) which weakens bonds, brings molecules closer together and puts them into a more favourable orientation.

2. Reaction takes place

3. Products are desorbed (leaves the surface)

Strength of adsorption:
Too strong (eg. V, Cr, Mn)

Reactants cannot move around surface

Products cannot desorb

Too Weak

Reactants not adsorbed (eg. Cu)

Ideal (eg. Fe, Co, Ni)
Nature of catalyst
Large Surface
Spread thinly over ceramic honeycomb
Poisoning
Soem substances may block active sites (I.e the adsorb and will not come off) - they can ruin a catalyst
Examples
S in Haber process
Pb in catalytic converters

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Copper(I) complexes have a full d-sub shell, meaning there are no empty orbitals to allow electrons to move between energy levels. Therefore, Cu(I) complexes don't absorb light in the visible range, appearing colourless. Similarly Sc³⁺ ions have an empty d-sub shells so there are no electrons to move between d-orbtials.

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What is the effect of splitting on electronic structure?

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What is the effect of splitting on electronic structure?

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To work out how the electrons are arranged in these orbitals, we must follow the same rules as for all other electronic structures. In the split orbitals, the lower three orbitals are filled first world 1 electron each, before the electrons pair up. Once these three orbitals are filled with 6 electrons, the higher 2 orbitals are filled.

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Colour in transition metal ions & complexes

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Colour in transition metal ions & complexes

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Transition metals with Ligands
In an octahedral complex, 6 negatively charged ligands approach the transition metal ion along the direction of the 3 axes.
The negative charges repel the electrons in the orbitals that point along the axes, making the orbital less stable.
This means that the orbitals are no longer degenerate (have the same energies.
The orbitals that don’t point along the axis aren’t affected.
The splitting of the d-orbitals gives 2 sets of orbitals close together in energy.
An electron in one of the d-orbitals can move from the lower to the upper set of orbitals by gaining energy absorbed in the form of light.
Only one frequency (colour) of light is absorbed, which corresponds to the energy gap between orbitals. The light that isn’t absorbed gives the complex its colour.
The d-d transitions depends on the amount of splitting between the d-orbitals and his varies between ions of different transition metal complexes. Therefore the frequencies of the light absorbed also differ, leading to different colours for different complexes. Different ligands give different splitting of orbitals.
Example
Compounds containing the complex [Cu(H₂O)₆]²⁺ are typically blue as they absorb all other colours. Compound containing [Fe(H₂O)₆)]³⁺ are yellow.

Transition metals without Ligands
Transition metals are only coloured in complexes. If there are no ligands around the metal ion, the compound would be colourless.

Transition metals with no ligands have 5 degenerate d-orbitals (5 d-ortbitals with the same energy.

The first 3 orbitals (dxz, dxy, dyz), point between a pair of axes

The last 2 orbitals (dx ²-y ², dz ²) point along the axes

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Copper complexes can show the variation in transition metal complexes. [Cu(H₂O)₆]²⁺, [Cu(NH₃)₄(H₂O)₂]²⁺ & [CuCl₄]²⁻ all have Cu²⁺ ions, but their different structure & properties is due to their different ligands. As well as cobalt forming [Co(H₂O)₆]²⁺ & [CoCl₄]²⁻ both containing Co2+ ions.

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Transition metal ions are small & can have large positive charges. They have many orbitals available for bonding, many are empty. Electron-rich molecules have lone pairs, so these can form coordinate bonds with the empty orbitals on the transition metals ions:

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Oxidation states for the transition elements 

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Oxidation states for the transition elements 

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Most common oxidation states to know:

Cr - +3, +5

Mn - +2, +3, +4, +6, +7

Fe - +2, +3

Co - +2, +3

Cu - +1, +2

They can form different oxidation states because the energies of the 4s & 3d-orbitals are very similar so the energy required to remove any of these electrons is similar. As the elements form compounds energy is released, either through the formation of covalent bonds or when the ionic lattice forms. The energy needed to reach higher oxidation states and the energy released in compound formation is finely balanced allowing a range of oxidation states to form.
The oxidation state favoured by each metals depends on many factors. The oxidising power of the other atom in the compound is one factor, so when iron metal reacts with chlorine gas it produces iron(III) chloride, when it reacts with iodine vapour the product is iron(III) oxide. The iodine is a much weaker oxidising agent than chlorine and so it cannot oxidise the iron atom the +3 oxidation state.

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When electrons are removed to form ions, the 4s electrons are lost first. This is because the 4s & 3d orbitals are very close together in energy, so on balance it is more energetically favourable to lose these 4s electrons before the 3d electrons.

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Electronic configuration for the elements 

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Electronic configuration for the elements 

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The 4s orbital is filled before 3d, even though it is written with 3d first.
2 transition elements don't obey this rule. Chromium & copper are different.

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D-block consists of elements from scandium to zinc. These elements are very different from the metals of the s-block. Transition metals are those with partially filled d-orbitals(scandium to nickel - partially filled d-orbitals in unreacted metals). Copper has full d-orbitals as a metal, but in most compounds has partially filled d-orbitals. Zinc has filled d-orbitals either way. Its d-sub shell is never partially filled, it's not a transition element.

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Tetrahedral ligands(4 bonds) less common

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Tetrahedral ligands(4 bonds) less common

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[CuCl₄]²⁻ - yellow or green complex

[CoCl₄]²⁻ - blue complex

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Ligand               Formula             Name
Chloride             Cl-                           Chloro
Cyanide              NC-                         Cyano
Hydroxide          HO-                        Hydroxo
Oxide                 O2-                            Oxo
Water                 H₂O                    Aqua
Ammonia          NH3                         Amine

They donate a lone pair of electron into vacant orbitals on the central species
Some Ligands attach themselves using two or more lone pairs
Classified by the number of lone pairs they use, not the number they have
Multidentate and bidentate ligands lead to more stable complexes

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[Fe(H₂O)₆]²⁺ - pale green complex

[Fe(H₂O)₆]³⁺ - yellow complex

[Cu(H₂O)₆]²⁺ - blue complex

[Cr(H₂O)₆]3+ - dark green complex

[Co(H₂O)₆]2+ - pink complex

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1. Variable Oxidation States

2. Coloured Ions

3. Form Complexes (Ligands form co-ordinated. Bonds to the metal ion)

4. Show Catalytic Activity

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The size of the energy gap between the d orbitals, and therefore the colour, are affected by changes in:

1. The metal eg. [Cu(H₂O)₆]²⁺ = Blue, [Fe(H₂O)₆]2+ = Green

2. The oxidation state eg. [Fe(H₂O)₆]2+ = Green, [Fe(H₂O)₆]3+ = Pale Violet

3. The ligands eg. [Cu(H₂O)₆]2+ = Blue, [Cu(H₂O)₂(NH₃)₄]2+ = Deep Blue

4. The co-ordination number  eg.  [Cu(H₂O)₆]2+ = Blue, [CuCl₄]²⁻ = Yellow

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Energy Levels of d orbitals in an Octahedral Field

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Energy Levels of d orbitals in an Octahedral Field

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The positive metal ion is attracted to the negative charges (Ligands) but electrons in the d orbitals are repelled by them

Although there is an overall attraction, the d orbitals no longer have the same energy

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V₂O5 - Catalyst in contact process

Mn2+ - Autocatalyst in MnO₄⁻/C₂O₄ 2- titrations

Haemoglobin - Contains Fe2+ which allows O₂ to bond and be carried to where needed.

[Pt(NH₃)₂Cl₂] - Cisplatin - anti cancer drug

[Ag(NH₃)₂]+ - Used in Tollen’s reagent to distinguish aldehydes and ketones

[Ag(CN)₂]⁻ - Used in silver electroplating

[Ag(S₂O₃)₂]³⁻ - Formed in photography to remove unreacted AgX from the film

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A lone pair acceptor

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A lone pair donor (Ligands are Lewis bases)

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The number of co-ordinate bonds from ligand(s) to metal ions.

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A metal ion with Ligands co-ordinately bonded to it.

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A small molecule with a lone pair that forms co-ordinate bonds to a transition metal.

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Zinc is not a transition metal because it’s d-subshell is complete.

4s subshell fills AND empties before 3d

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Metals that contain an incomplete d subshell  in atoms or ions.