Topic 15: Transition Metals

Transition metals

Transition metals are d-block elements that can form one or more stable ions with incompletely filled d-orbitals.

d-block elements that are not transition metals

Scandium and zinc.

How transition metals form ions

Transition metals lose electrons from their 4s orbital first, then their 3d orbitals, as the 4s subshell has the highest energy.

Electron configuration of Cu²⁺

1s²2s²2p⁶3s²3p⁶3d⁹

Electron configuration of Cr⁶⁺

1s²2s²2p⁶3s²3p⁶

Why transition metals show variable oxidation number

• The 3d and 4s subshells are very close in energy.

• This allows transition metals to first lose electrons from their 4s subshell and then also lose more electrons from the d-subshell without a large increase in energy required, so the successive ionisation energies increase gradually.

Ligand

A ligand is a species that uses a lone pair of electrons to form a dative covalent bond with a central metal ion.

Complex

• A complex is a central metal ion surrounded by ligands, bonded to the metal ion with dative bonds.

• A complex with a charge is a complex ion.

Coordination number

The number of dative covalent bonds in a complex is its coordination number.

Why the solutions of aqueous ions of transition metals are coloured

• The d-subshell of the ions of transition metals are incompletely filled, so when water ligands attach to the ion, the d-orbitals split into two different energy levels.

• The electrons in the d-orbitals absorb energy from the visible spectrum, causing them to move from a lower energy d-orbital to a higher energy d-orbital.

• The remaining colour of light that is not absorbed is transmitted and seen as colour.

Why zinc ions do not form coloured solutions

• Zn²⁺ ions have a full 3d subshell.

• This means electrons cannot move between 3d orbitals.

Why aluminium ions do not form coloured solutions

• Al³⁺ ions have no electrons in their 3d subshell.

• This means there are no d-orbital electrons that can absorb energy and move between 3d orbitals.

Shapes of different complexes

• When the coordination number is 6, the complex has an octahedral shape.

• When the coordination number is 4, the complex has a tetrahedral shape or sometimes a square planar shape.

• When the coordination number is 2, the complex has a linear shape.

Monodentate ligands

• A ligand that forms one dative bond with a metal ion.

• Examples are H₂O, OH^-, and NH₃, as they all only have one atom with a lone pair of electrons.

Ligands that form octahedral complexes

NH₃, OH^-, and NH₃, as they are relatively small and can fit around the central ion.

Ligands that form tetrahedral complexes

Cl^- ions, as they are relatively large, so more than four cannot fit around the central ion.

Complex ion of Fe²⁺ with H₂O

• [Fe(H₂O)₆]²⁺

• Coordination number is 6.

Complex ion of Cu²⁺ with Cl^-

• [CuCl₄]^2-

• Coordination number is 4 and is tetrahedral.

Cis-platin complex

• [Pt(NH₃)₂Cl₂]

• Consists of Pt²⁺ ion, two ammonia ligands, and two Cl^- ion ligands.

• The cis- prefix indicates that identical ligands are next to each other.

• The complex has a square planar shape with bond angles of 90^o.

Usage of cis-platin

• Cis-platin is used to treat cancer.

• When cis-platin enters the cell, it loses its chloride ligands, which are replaced by water molecules.

• Cis-platin becomes a highly reactive, positively charged ion.

• The water ligands are removed and replaced by dative covalent bonds to two adjacent guanine bases on the same strand of DNA, which is made possible by the two water ligands being on the same side of the complex.

• This disrupts the DNA structure and stops the cancer cell from dividing.

• Trans-platin is not able to be used as cancer treatment because the water ligands are on opposite sides of the complex, so they cannot form dative covalent bonds with adjacent guanine bases on the same strand of DNA.

Bidentate ligands

• A ligand that forms two dative bonds with a metal ion.

• An example is NH₂CH₂CH₂NH₂, which uses the lone pair of electrons on each nitrogen atom to form two dative bonds with the metal ion. It is sometimes abbreviated as ‘en’.

• Three NH₂CH₂CH₂NH₂ ligands bond with a metal ion to form an octahedral shape and a complex ion with a coordination number of 6.

Multidentate ligands

• A ligand that forms several dative bonds with a metal ion.

• An example is EDTA^4-, which is formed when EDTA’s COOH groups each lose a H⁺ ion. It can form 6 dative bonds with a metal ion, four with its lone electron pairs on the negative oxygen of its COO^- groups and 2 with the lone electron pairs on its nitrogen atoms.

• One EDTA^4- ligand bonds with a metal ion to form an octahedral shape and a complex ion with a coordination number of 6.

Substitution of a monodentate ligand by a bidentate or multidentate ligand

• The substitution of a monodentate ligand by a bidentate or multidentate ligand creates a more stable complex ion because the reaction has a large positive ΔS_system, meaning the system goes from a less stable state to a more stable state.

• For example, [Cu(H₂O)₆]²⁺ + EDTA^4- → [Cu(EDTA)]^2- + 6H₂O, shows that ΔS_system is positive, as 2 moles of reactants forms 7 moles of products, meaning that there is an increase in disorder.

Haemoglobin

• Haemoglobin is an iron (II) complex containing a multidentate ligand and 4 Fe²⁺ ions.

• The largest part of haemoglobin is the protein, globin, which contains four haem groups.

• Each haem group contains 4 nitrogen atoms, and each haem group forms four dative bonds, using their four nitrogen atoms, to each Fe²⁺ ion.

• Another dative bond forms between the globin and each Fe²⁺ ion.

• When haemoglobin carries oxygen, the oxygen molecule acts as a ligand, using one of its lone pairs of electrons to form a dative bond with the Fe²⁺ ion.

Carbon monoxide and haemoglobin

• Carbon monoxide acts as a ligand by using its carbon atom’s lone pair of electrons to form a dative bond with a metal ion.

• The strength of the dative bond between Fe²⁺ and carbon monoxide is a lot stronger than the strength of the dative bond between oxygen and Fe²⁺, so carbon monoxide readily replaces the oxygen from haemoglobin, causing a ligand exchange reaction. 

• This dative bond is so strong that the reaction is irreversible.

Causes of colour changes in transition metal ions

• A change in oxidation number of the ion may change the colour of the ion.

  • When a metal ion has a higher oxidation number, there is a stronger electrostatic attraction to the ligands, causing the d-orbitals to be split further apart, creating a larger energy gap.

• A change in ligand of the ion may change the colour of the complex ion.

  • The greater the charge density of a ligand, the greater the electrostatic attraction to the metal ion, causing the d-orbitals to be split further apart, creating a larger energy gap.

• A change in the coordination number of the complex ion may change the colour of the complex ion.

  • A higher coordination number of the same ligands will cause the d-orbitals to be split further apart, causing a larger energy gap.

Reaction of Cu²⁺ (aq) with aqueous sodium hydroxide

•  [Cu(H₂O)₆]²⁺ + 2OH^- → [Cu(H₂O)₄(OH)₂] + 2H₂O

• A blue solution forms a blue precipitate.

• Two of the water ligands transfer a H⁺ ion to each OH^- ion, so this is an acid-base reaction.

• There is no change if excess NaOH (aq) is added.

Reaction of Cu²⁺ (aq) with aqueous ammonia

• [Cu(H₂O)₆]²⁺ + 2NH₃ → [Cu(H₂O)₄(OH)₂] + 2NH₄⁺

• Two of the water ligands transfer a H⁺ ion to each ammonia molecule, so this is an acid-base reaction.

• A blue solution forms a blue precipitate

• If excess aqueous ammonia is added, the precipitate dissolves and the solution turns dark-blue: [Cu(H₂O)₄(OH)₂] + 4NH₃ → [Cu(NH₃)₄(H₂O)₂]²⁺ + 2H₂O + 2OH^-, which is a ligand exchange reaction.

• The overall equation when excess aqueous ammonia is added is [Cu(H₂O)₆]²⁺ + 4NH₃ → [Cu(NH₃)₄(H₂O)₂]²⁺ + 4H₂O

Reaction of Fe²⁺ (aq) with aqueous sodium hydroxide

• [Fe(H₂O)₆]²⁺ + 2OH^- → [Fe(H₂O)₄(OH)₂] + 2H₂O

• Two of the water ligands transfer a H⁺ ion to each OH^- ion, so this is an acid-base reaction.

• A green solution forms a green precipitate.

• There is no change if excess NaOH (aq) is added.

Reaction of Fe²⁺ (aq) with aqueous ammonia

• [Fe(H₂O)₆]²⁺ + 2NH₃ → [Fe(H₂O)₄(OH)₂] + 2NH₄⁺

• Two of the water ligands transfer a H⁺ ion to each ammonia molecule, so this is an acid-base reaction.

• A green solution forms a green precipitate.

• There is no change if excess aqueous ammonia is added.

Reaction of Fe³⁺ (aq) with aqueous sodium hydroxide

• [Fe(H₂O)₆]³⁺ + 3OH^- → [Fe(H₂O)₃(OH)₃] + 3H₂O

• Three of the water ligands transfer a H⁺ ion to each OH^- ion, so this is an acid-base reaction.

• A yellow solution forms a brown precipitate.

• There is no change if excess NaOH (aq) is added.

Reaction of Fe³⁺ (aq) with aqueous ammonia

• [Fe(H₂O)₆]³⁺ + 3NH₃ → [Fe(H₂O)₃(OH)₃] + 3NH₄⁺

• Three of the water ligands transfer a H⁺ ion to each ammonia molecule, so this is an acid-base reaction.

• A yellow-brown solution forms a brown precipitate.

• There is no change if excess aqueous ammonia is added.

Reaction of Co²⁺ (aq) with aqueous sodium hydroxide

• [Co(H₂O)₆]²⁺ + 2OH^- → [Co(H₂O)₄(OH)₂] + 2H₂O

• Two of the water ligands transfer a H⁺ ion to each OH^- ion, so this is an acid-base reaction.

• A pink solution forms a blue precipitate.

• There is no change if excess NaOH (aq) is added.

Reaction of Co²⁺ (aq) with aqueous ammonia

• [Co(H₂O)₆]²⁺ + 2NH₃ → [Co(H₂O)₄(OH)₂] + 2NH₄⁺

• Two of the water ligands transfer a H⁺ ion to each ammonia molecule, so this is an acid-base reaction.

• A pink solution forms a blue precipitate.

• If excess aqueous ammonia is added, the precipitate dissolves and forms a yellow solution: [Co(H₂O)₄(OH)₂] + 6NH₃ → [Co(NH₃)₆]²⁺ + 4H₂O + 2OH^-, which is a ligand exchange reaction.

• The overall equation when excess aqueous ammonia is added is [Co(H₂O)₆]²⁺ + 6NH₃ → [Co(NH₃)₆]²⁺ + 6H₂O

Reaction of Cr³⁺ (aq) with aqueous sodium hydroxide

• [Cr(H₂O)₆]³⁺ + 3OH^- → [Cr(H₂O)₃(OH)₃] + 3H₂O

• Three of the water ligands transfer a H⁺ ion to each OH^- ion, so this is an acid-base reaction.

• A green solution forms a green precipitate.

• If excess aqueous sodium hydroxide is added, the precipitate dissolves to form a green solution: [Cr(H₂O)₃(OH)₃] + 3OH^- → [Cr(OH)₆]^3- + 3H₂O

• Three of the water ligands transfer a further three H⁺ ions to the OH^- ions, so this is an acid-base reaction.

Reaction of Cr³⁺ (aq) with aqueous ammonia

• [Cr(H₂O)₆]³⁺ + 3NH₃ → [Cr(H₂O)₃(OH)₃] + 3NH₄⁺

• Three of the water ligands transfer a H⁺ ion to each ammonia molecule, so this is an acid-base reaction.

• A green solution forms a green precipitate.

• If excess aqueous ammonia is added, the precipitate dissolves to form a violet solution: [Cr(H₂O)₃(OH)₃] + 6NH₃  → [Cr(NH₃)₆]³⁺ + 3H₂O + 3OH^-, which is a ligand exchange reaction.

• The overall equation when excess aqueous ammonia is added is [Cr(H₂O)₆]³⁺ + 6NH₃ → [Cr(NH₃)₆]³⁺ + 6H₂O

Formation of [CuCl₄]^2- from [Cu(H₂O)₆]²⁺

• [Cu(H₂O)₆]²⁺ + 4Cl^- → [CuCl₄]^2- + 6H₂O

• This is a ligand exchange reaction.

• The colour of the solution changes from blue to yellow.

Formation of [CoCl₄]^2- from [Co(H₂O)₆]²⁺

• [Co(H₂O)₆]²⁺ + 4Cl^- → [CuCl₄]^2- + 6H₂O

• This is a ligand exchange reaction.

• The colour of the solution changes from pink to blue.

Substitution of small, uncharged ligands by larger, charged ligands

This ligand exchange reaction often results in a change in coordination number, as less of the larger, charged ligands can fit around the central metal ion, as there is more electrostatic repulsion between them and they are large.

Amphoteric behaviour of transition metal hydroxides with chromium (III) hydroxide as an example

• Transition metal hydroxides can both accept H⁺ ions and donate them, so they are amphoteric.

• [Cr(H₂O)₃(OH)₃] + 3H⁺ → [Cr(H₂O)₆]³⁺

• [Cr(H₂O)₃(OH)₃] + 3OH^- → [Cr(OH)₆]^3- + 3H₂O

Equilibrium between chromate (VI) and dichromate (VI) ions

• 2CrO₄^2- + 2H⁺ ⇌ Cr₂O₇^2- + H₂O

• When the concentration of H⁺ increases and the solution becomes more acidic, C₂O₇^2- forms, causing an orange colour.

• When the concentration of OH^- increases and the solution becomes more basic, 2CrO₄^2- forms, causing a yellow colour.

Reaction of dichromate (VI) ions with zinc

• In acidic conditions, zinc reduces dichromate (VI) ions first to Cr³⁺ ions and then to Cr²⁺ ions.

• This is made possible because of the E^⦵ values of the relevant reactions.

Reaction of Cr³⁺ ions with hydrogen peroxide

• In alkaline conditions, hydrogen peroxide (H₂O₂) oxidises Cr³⁺ ions to dichromate (VI) ions.

• This is made possible because of the E^⦵ values of the relevant reactions.

Vanadium oxidation state colours of solutions

• +2: purple

• +3: green

• +4: blue

• +5: yellow

Heterogenous catalyst

• A catalyst in a different state to the reactants.

• They are often solids used in reactions involving gases, providing a surface for the reactant gas molecules to adsorb onto. The reactant molecules adsorb onto the active sites of the catalyst. Bonds between the catalyst and the reactant molecules are formed, weakening the bonds between the atoms of the reactant molecules, allowing a reaction to take place between the reactant molecules on the surface of the catalyst. The product then desorbs from the surface and more reactant molecules take their place.

Homogenous catalyst

• A catalyst in the same state as the reactants.

• The catalysed reaction proceeds via an intermediate species.

Usage of transition metals as catalysts

• Transition metals and their compounds can act as both heterogeneous and homogeneous catalysts.

• Transition metals and their compounds can act as catalysts because they have variable oxidation states, allowing them to change their oxidation state back and forth as they react with the reactants of the reaction.

How vanadium (V) oxide is used as a catalyst in the contact process

• The contact process produces sulfuric acid.

• The reaction of sulfur dioxide with oxygen to form sulfur trioxide is catalysed by vanadium (V) oxide: 2SO₂ + O₂ ⇌ 2SO₃.

• Vanadium (V) oxide is reduced by sulfur dioxide into vanadium (IV) oxide to form sulfur trioxide: SO₂ + V₂O₅ → V₂O₄ + SO₃

• Vanadium (IV) oxide is the intermediate species.

• The vanadium (IV) oxide then reacts when oxygen to reform vanadium (V) oxide, allowing it to fulfill its role as a catalyst: 2V₂O₄ + O₂ → 2V₂O₅

Catalytic converter

• A catalytic converter reduces carbon monoxide and nitrogen monoxide emissions from combustion engines.

• The CO and NO molecules adsorb onto the surface of the catalyst.

• The catalyst weakens the bonds between the atoms of the molecules, allowing them to react: 2CO + 2NO → 2CO₂ + N₂

• The nitrogen and carbon dioxide products then desorb from the surface of the catalyst so more reactant molecules can take their place.

Reaction between I^- and S₂O₈^2- ions

• S₂O₈^2- ions oxidise I^- ions: 2I^- + S₂O₈^2- → I₂ + 2SO₄^2-

• The reaction is slow at room temperature because the ions repel each other.

• The reaction is catalysed by Fe²⁺ ions.

• All the reactants and the catalyst are in the aqueous phase, so it is homogeneous catalysis.

• When catalysed, Fe²⁺ reduce S₂O₈^2- ions, which are attracted to each other, to produce the intermediate species Fe³⁺: S₂O₈^2- + 2Fe²⁺ → 2SO₄^2- + 2Fe³⁺

• The Fe³⁺ ions oxidise I^- ions, which are attracted to each other: 2Fe³⁺ + 2I^- → 2Fe²⁺ + I₂

• This reforms the Fe³⁺, allowing it to fulfill its role as a catalyst.

Reaction between manganate (VII) and ethanedioate ions

• Manganate (VII) ions oxidise ethanedioate ions: 2MnO₄^- + 5C₂O₄^2- → 2Mn²⁺ + 5CO₂ + 8H₂O

• The initial rate of reaction is low because the ions repel each other.

• Mn²⁺ autocatalyses the reaction, so the rate of reaction increases as the reaction proceeds.

• Mn²⁺ is oxidised to Mn³⁺, which is the intermediate species, when it reacts with MnO₄^-: Mn²⁺ + MnO₄^- + 8H⁺ → 5Mn³⁺ + 4H₂O

• The Mn³⁺ ions are then reduced back to Mn²⁺ when they react with C₂O₄^2-: 2Mn³⁺ + C₂O₄^2- → 2CO₂ + 2Mn²⁺

• This reforms the Mn²⁺, allowing it to fulfill its role as a catalyst.