General Characteristic Chemical Properties of the First Set of Transition Elements, Titanium to Copper

Specification Reference Inorganic Chemistry, Chemistry of transition elements 28.2

Quick Notes

Ligands: Species with a lone pair of electrons that form coordinate (dative covalent) bonds to transition metal ions.
Types of Ligands:
- Monodentate: can form one bond (e.g., H₂O, NH₃, Cl⁻, CN⁻).
- Bidentate: can form two bonds (e.g., ethane-1,2-diamine "en", C₂O₄²⁻).
- Polydentate: can form multiple bonds (e.g., EDTA⁴⁻).
Complex Ions: Transition metals form complex ions surrounded by ligands. Shapes:
- Linear (180°)
- Tetrahedral (~109.5°)
- Square Planar (90°)
- Octahedral (90°)
Coordination Number: Number of coordinate bonds to the central metal ion.
Ligand Exchange: Ligands can be swapped in complexes (partial or full substitution).
Colour Formation: d-orbitals split when ligands bond, allowing visible light absorption and colour.
Redox Reactions: Transition elements easily change oxidation state. Predict feasibility using E° values:
- Higher E° = easier reduction (strong oxidising agent).
- Lower E° = easier oxidation (strong reducing agent).
Common redox titrations:

Full Notes

Complexes

Transition metals form complex ions by accepting lone pairs from ligands via dative covalent bonds.

Water acting as a ligand via a lone pair to a central metal ion.

For example
Water molecules (H₂O) are able to act as ligands as the oxygen atom can use one its lone pairs electrons to form a co-ordinate bond to a central metal atom or ion.

Ligands

A ligand is a molecule or ion with a lone pair of electrons that can form a coordinate bond to a central metal ion.

The bond is a dative covalent bond – both electrons in the bond come from the ligand.

Types of Ligands

We classify ligands based on the number of co-ordinate bonds they can form. (see ligand substitution for more detail).

Monodentate: Donates one pair of electrons (e.g., H₂O, NH₃, Cl⁻).
Bidentate: Donates two pairs of electrons (e.g., ethane-1,2-diamine).
Multidentate: Donates multiple pairs of electrons (e.g., EDTA⁴⁻).

Complex Ions

A complex ion is a species formed when ligands bond to a central metal ion.

Example: [Cu(H₂O)₆]²⁺ is a copper ion surrounded by six water ligands.

The formulas of complex ions are written in square brackets with the overall charge of the complex ion shown as a superscript.

How to write complex ion formulae with charges.

Coordination Number

The co-ordination number of a complex is the number of coordinate bonds formed with the central metal ion and determines the geometry (shape) of the complex.

Linear, tetrahedral, square planar and octahedral complex shapes.

Matt’s exam tip

Don’t forget the co-ordination number of a complex and the number of ligands don’t have to be the same.

The shape of a complex is determined by the co-ordination number and type of ligands bonded to a central metal atom or ion. The most common shapes are octahedral, tetrahedral, square planar and linear.

Shapes and Bond Angles of Complexes

Shape	Co-ordination number	Typical ligands	Common isomerism	Example complex
Octahedral	6	H₂O, NH₃ (small, neutral)	Cis–trans; Optical (with bidentate)	[Cu(H₂O)₆]²⁺, [Cr(NH₃)₆]³⁺
Tetrahedral	4	Cl⁻ (larger anion)	Usually none	[CuCl₄]²⁻
Square planar	4	Pt(II), Ni(II) complexes	Cis–trans	[Pt(NH₃)₂Cl₂] (cisplatin)
Linear	2	NH₃ (ammine)	None	[Ag(NH₃)₂]⁺ (Tollens’ reagent)

Ligand Exchange Reactions

Ligands in a complex can be partially or fully replaced by other ligands.

Example [Cu(H₂O)₆]²⁺ with Cl⁻ and NH₃

Ligand exchange of copper(II) aqua complex with chloride and ammonia.

Addition of Cl- forms yellow solution [CuCl₄]²⁻, changing shape to tetrahedral.
Addition of Excess NH₃ forms deep deep blue solution of [Cu(NH₃)₄(H₂O)₂]²⁺.

Example[Co(H₂O)₆]2⁺ with Cl⁻ and NH₃

Ligand exchange of cobalt(II) aqua complex with chloride and ammonia.

Addition of Cl- forms blue solution [CoCl₄]²⁻, changing shape to tetrahedral.
Addition of Excess NH₃ forms deep yellow solution of [Co(NH₃)6]²⁺.

Key Point - if only a limited amount of NH3(aq) or OH-(aq) is added, a metal hydroxide precipitate forms. This isn’t technically ligand substitution as the complex ion is acting as an acid and water ligands are losing H+(aq) ions, become OH- ligands.

Formation of metal hydroxide precipitates with limited NH3/OH-.

Matt’s exam tip

Remember copper (II) undergoes only partial substitution with NH3 forming [Cu(NH₃)₄(H₂O)₂]²⁺.

Redox Reactions and Electrode Potentials (E°)

Note redox potentials and electrode potentials have been covered in more detail here.

Transition metals can change oxidation state, meaning they are often involved in redox reactions.
Standard electrode potentials (E°) help predict feasibility (for more on this see electrode potentials).
- If E°(reduction) is more positive, it’s more likely to happen.

Redox Reactions You Should Know

MnO₄⁻ and C₂O₄²⁻ in Acidic Solution

Permanganate–oxalate redox titration in acid.

Used in redox titrations, purple MnO₄⁻ becomes colourless Mn²⁺.

MnO₄⁻ and Fe²⁺ in Acidic Solution

Permanganate–iron(II) redox titration in acid.

Again, purple to colourless, self-indicating end-point.

Cu²⁺ and I⁻

Copper(II) with iodide forming CuI and iodine.

Note: CuI is white ppt and I₂ gives a brown solution

Calculations in Redox Titrations

You may be asked to:

Use moles, volumes, and concentrations to find unknown quantities.
Apply stoichiometry from the redox equations.

Worked Example – Finding the % of Iron in an Iron Tablet

Problem:
An iron tablet was dissolved and made up to 250.0 cm³. 25.0 cm³ of this solution was titrated with 0.0200 mol dm⁻³ KMnO₄. The average titre was 23.60 cm³. The tablet’s mass was 2.50 g (larger than before). Calculate the percentage by mass of iron (Fe²⁺) in the tablet. (Relative atomic mass of Fe = 55.8)

Step 1: Write the redox equation
MnO₄⁻ + 5Fe²⁺ + 8H⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O
Step 2: Calculate moles of MnO₄⁻ used
Moles = concentration × volume (in dm³)
Moles MnO₄⁻ = 0.0200 × (23.60 ÷ 1000) = 4.72 × 10⁻⁴ mol
Step 3: Find moles of Fe²⁺
From the stoichiometry, 1 mol MnO₄⁻ reacts with 5 mol Fe²⁺.
Thus: Moles Fe²⁺ = 5 × 4.72 × 10⁻⁴ = 2.36 × 10⁻³ mol (in 25.0 cm³)
Step 4: Scale up to 250.0 cm³
The whole solution is 10 times larger than the sample.
Total moles Fe²⁺ = 2.36 × 10⁻³ × 10 = 2.36 × 10⁻² mol
Step 5: Calculate mass of Fe
mass = moles × Mr
mass of Fe = 2.36 × 10⁻² × 55.8 = 1.317 g
Step 6: Find % of iron in the tablet
% Fe = (mass of Fe ÷ mass of tablet) × 100
% Fe = (1.317 ÷ 2.50) × 100 = 52.7%

Matt’s exam tip

Look out for dilutions with redox style titration questions. In this example, we have to remember the whole solution is 10x larger than the sample used in the titration. This is very common in these kind of exam questions.

Summary

Ligands: Species with a lone pair of electrons that form coordinate (dative covalent) bonds to transition metal ions.
Types of ligands: Monodentate, Bidentate, Polydentate (e.g., EDTA⁴⁻).
Coordination number = number of coordinate bonds.
Ligands can be swapped in complexes (partial or full substitution) in ligand exchange.
Colour formation:
d-orbitals split and electrons absorb visible light to become excited to higher energy, the observed colour depends on absorbed light.
Redox reactions: Transition elements easily change oxidation state and we can use E° values to predict feasibility.