Chemistry (9620) | Transition metals

Hello, Future Chemist! Welcome to Transition Metals!

This chapter is all about the "rainbow section" of the Periodic Table—the Transition Metals (the d-block elements, specifically from Titanium (Ti) to Copper (Cu) in Period 4). If you've ever wondered why copper wiring is green when corroded, or why some paints are so vividly coloured, this is where you find the answers!

Transition metals are incredibly important in industry and biology (think iron in your blood!). Don't worry if the names of complexes look complicated; we will break down all the rules clearly. Let's dive into their unique world!

3.2.5 General Properties of Transition Metals

Transition metals are defined by their ability to form at least one ion that has an incomplete d sub-level (in the syllabus, this is limited to elements Ti to Cu). This incomplete d-shell is the secret behind all their fascinating chemical characteristics.

Key Characteristics (The Four Cs)

Transition metals exhibit four major characteristics that set them apart from s-block metals (like Na or Mg).

• Complex Formation: They readily form coordination complexes.
• Coloured Ions: Their compounds are often brilliantly coloured.
• Catalytic Activity: They (or their compounds) act as powerful catalysts.
• Variable Oxidation States: They can exist in many different stable oxidation states.

Complex Formation and Ligands

When a transition metal ion sits in a solution, it acts as a central hub, attracting other molecules or ions around it.

What is a Ligand and a Complex?

• Ligand: A molecule or ion that can donate a pair of electrons (a lone pair) to a central transition metal ion to form a co-ordinate bond (dative covalent bond). Ligands are therefore Lewis bases. Examples: H₂O, NH₃, Cl⁻.
• Complex: The structure formed when a central metal ion is surrounded by ligands. The whole complex is often enclosed in square brackets. Example: \([Cu(H_2O)_6]^{2+}\).
• Co-ordination Number (CN): This is the number of co-ordinate bonds formed between the central metal ion and the ligands. The most common CNs are 6, 4, and 2.

Types of Ligands

We classify ligands based on how many lone pairs they "bite" onto the central metal ion:

1. Monodentate Ligands: Form one co-ordinate bond. They have one donor atom.
   Examples in the syllabus: H₂O, NH₃, Cl⁻.
2. Bidentate Ligands: Form two co-ordinate bonds (two "bites").
   Examples in the syllabus: Ethan-1,2-diamine (\(H_2NCH_2CH_2NH_2\)) and Ethandioate ion (\(C_2O_4^{2-}\)).
3. Multidentate Ligands: Form multiple co-ordinate bonds (more than two).
   Example in the syllabus: EDTA⁴⁻ (ethylenediaminetetraacetate ion), which is hexadentate (forms 6 bonds).

Ligand Substitution Reactions

Ligands can be swapped out for new ligands in a reaction.

• H₂O and NH₃ Exchange: Since H₂O and NH₃ are similar in size and both uncharged, substitution often occurs without a change in co-ordination number (usually stays at CN=6).
  Example: \([Co(H_2O)_6]^{2+} + 6NH_3 \rightleftharpoons [Co(NH_3)_6]^{2+} + 6H_2O\).
• Cl⁻ Exchange: The chloride ion (\(Cl^-\)) is much larger than H₂O or NH₃. When \(Cl^-\) substitutes H₂O, there is often a change in co-ordination number (usually from 6 to 4).
  Example: \([Cu(H_2O)_6]^{2+} + 4Cl^- \rightleftharpoons [CuCl_4]^{2-} + 6H_2O\).

The Chelate Effect (The "Backpack" Analogy)

The replacement of monodentate ligands by bi- or multidentate ligands is thermodynamically favourable. This preference for polydentate ligands is called the chelate effect.

Don't worry if this seems tricky! The explanation rests on entropy (\(\Delta S\)).

1. Imagine replacing six water molecules (monodentate) with three ethan-1,2-diamine molecules (bidentate).

2. Enthalpy (\(\Delta H\)): The six new bonds formed are chemically similar to the six old bonds broken. Thus, the enthalpy change (\(\Delta H\)) is usually very small (close to zero).

3. Entropy (\(\Delta S\)): You start with 4 particles (1 complex ion + 3 bidentate ligands) and end up with 7 particles (1 new complex + 6 displaced water ligands).
   Since the number of particles increases, the disorder (entropy, \(\Delta S\)) of the system increases significantly (\(\Delta S\) is positive).

4. Feasibility (\(\Delta G\)): The Gibbs Free Energy equation is \(\Delta G = \Delta H - T\Delta S\).
   Since \(\Delta H \approx 0\) and \(\Delta S\) is positive, \(\Delta G\) becomes negative. A negative \(\Delta G\) means the reaction is feasible and spontaneous.

Analogy: Replacing six small ligands with three large ligands is like switching six small shopping bags for three large backpacks. The backpacks (chelating ligands) hold the metal ion much more tightly and securely, and the process is driven by the "mess" created (the six water molecules floating off).

Shapes of Complex Ions and Isomerism

The shape of a complex ion depends mainly on the co-ordination number and the size of the ligands.

Common Shapes

• Octahedral (CN=6): Most common shape, especially with small ligands like H₂O and NH₃. The metal ion is at the center, surrounded by six ligands at 90° angles.
• Tetrahedral (CN=4): Common with larger ligands, such as \(Cl^-\). The bond angle is approximately 109.5°. Example: \([CuCl_4]^{2-}\).
• Square Planar (CN=4): A less common shape, often found in complexes of d⁸ ions (like Pt(II)).
• Linear (CN=2): Found with very small co-ordination numbers. Example: \([Ag(NH_3)_2]^+\) (used in Tollens' reagent).

Isomerism in Complexes

Isomers are compounds with the same formula but different arrangements of atoms in space.

1. Cis-Trans Isomerism (Geometrical Isomerism):
   • Occurs in octahedral and square planar complexes.
   • Requires at least two pairs of identical ligands.
   • Cis: Identical ligands are next to each other (at 90°).
   • Trans: Identical ligands are opposite each other (at 180°).

   Crucial Example: Cisplatin (cis-\([PtCl_2(NH_3)_2]\)) is the active anticancer drug (it binds to DNA); the trans isomer is inactive.

2. Optical Isomerism (Enantiomerism):
   • Occurs in octahedral complexes containing bidentate ligands.
   • The two isomers are non-superimposable mirror images of each other.

Key Takeaway for Shapes: Small ligands (H₂O, NH₃) favour Octahedral (CN=6). Large ligands (Cl⁻) favour Tetrahedral (CN=4). If you see "cis-trans" or "optical", think Octahedral complexes with bidentate ligands!

3.2.5.4 Formation of Coloured Ions (The Rainbow Effect)

One of the most characteristic properties of transition metals is the vibrant colour of their compounds.

Why are Transition Metal Complexes Coloured?

Colour arises from the movement of electrons within the incomplete d-sub-level.

1. When ligands approach the central metal ion, they affect the energy of the d-orbitals, causing them to split into two distinct energy levels.
2. An electron in the lower energy level (the ground state) can absorb a specific amount of energy (\(\Delta E\)) from visible light and jump up to the higher energy level (the excited state).
3. The relationship between the energy absorbed (\(\Delta E\)) and the wavelength (\(\lambda\)) of the light is given by:
   $ \Delta E = h\nu = \frac{hc}{\lambda} $
   where \(h\) is Planck's constant, \(c\) is the speed of light, and \(\nu\) is frequency.
4. The complex absorbs one colour (one wavelength) and transmits or reflects the remaining light. The colour we see is the mixture of the wavelengths that were not absorbed (the complementary colour).

Example: If a copper complex absorbs yellow light, it will appear blue (the complementary colour).

Factors Affecting Colour

Since colour depends directly on the energy difference (\(\Delta E\)) between the split d-orbitals, anything that changes \(\Delta E\) will change the colour:

• Oxidation State: \([Fe(H_2O)_6]^{2+}\) (pale green) vs. \([Fe(H_2O)_6]^{3+}\) (yellow/brown).
• Co-ordination Number: \([Cu(H_2O)_6]^{2+}\) (blue, CN=6) vs. \([CuCl_4]^{2-}\) (yellow/green, CN=4).
• Ligand Type: \([Ni(H_2O)_6]^{2+}\) (green) vs. \([Ni(NH_3)_6]^{2+}\) (purple).

Using Colour: Colorimetry

A colorimeter is an instrument used to measure the concentration of a coloured ion in a solution. It works by measuring how much light (of the colour absorbed by the ion) is blocked or absorbed by the sample.

Key Takeaway for Colour: Transition metal compounds are coloured because their d-electrons absorb specific wavelengths of visible light to jump to a higher energy level. The colour observed is the light that passes through.

3.2.5.5 Variable Oxidation States & Redox

Transition metals can exist in multiple stable oxidation states. This is another consequence of the d-electrons being close in energy, meaning different numbers of electrons can be easily removed or added.

Vanadium (A Great Example)

Vanadium is excellent for showing how oxidation states change colour. Starting with the Vanadate(V) ion, it can be progressively reduced by zinc metal in acidic solution:

1. Vanadate(V) ion: \(VO_2^+\) (aq) – Yellow (O.S. +5)
2. Reduced to Vanadium(IV) ion: \(VO^{2+}\) (aq) – Blue (O.S. +4)
3. Reduced to Vanadium(III) ion: \(V^{3+}\) (aq) – Green (O.S. +3)
4. Reduced to Vanadium(II) ion: \(V^{2+}\) (aq) – Violet/Lilac (O.S. +2)

Tip: You need to know these four colours and their corresponding oxidation states!

Redox Potential

The ability of a transition metal ion to change its oxidation state (its redox potential) is influenced by the surrounding ligands and the pH of the solution.

Redox Titrations

Because transition metal ions have distinct, deep colours in different oxidation states, they are used extensively in quantitative analysis (titrations).

• Permanganate Titration: Permanganate ions (\(MnO_4^-\), O.S. +7, purple) are powerful oxidising agents. They are used to titrate reducing agents like \(Fe^{2+}\) (O.S. +2) and ethanedioate ions (\(C_2O_4^{2-}\)).
• The end-point is signalled by the solution turning permanently pale pink due to the first drop of unreacted purple \(MnO_4^-\).
  The \(Mn^{2+}\) product formed in acidic solution is colourless (O.S. +2).

Tollens' Reagent: The reduction of the complex ion \([Ag(NH_3)_2]^+\) (a silver(I) complex) to metallic silver is used to distinguish aldehydes from ketones. This is a crucial redox reaction you must remember!

3.2.5.6 Catalysts (The Workhorses of Industry)

Transition metals and their compounds are excellent catalysts because they can easily change their oxidation states and form temporary bonds with reactants.

Heterogeneous Catalysis

The catalyst is in a different phase from the reactants (usually a solid catalyst and gaseous/liquid reactants). The reaction happens on the solid surface (active sites).

• Iron (Fe) in the Haber Process: Fe (solid) catalyses \(N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)\).
• Vanadium(V) Oxide (\(V_2O_5\)) in the Contact Process: \(V_2O_5\) (solid) catalyses the crucial step \(2SO_2 + O_2 \rightarrow 2SO_3\).
  Mechanism (involves changing oxidation states):

  \(2V_2O_5 + 2SO_2 \rightarrow 2V_2O_4 + 2SO_3\) (V(V) is reduced to V(IV))
  \(2V_2O_4 + O_2 \rightarrow 2V_2O_5\) (V(IV) is re-oxidised back to V(V))

  The \(V_2O_5\) is regenerated, showing its role as a catalyst.

Homogeneous Catalysis

The catalyst is in the same phase as the reactants (usually all aqueous ions). The reaction proceeds through an intermediate species.

• Fe²⁺/S₂O₈²⁻ Reaction: Iron(II) ions catalyse the reaction between iodide ions (\(I^-\)) and peroxodisulfate ions (\(S_2O_8^{2-}\)).
  Mechanism: Fe²⁺ acts as a stepping stone between the two reactants, changing oxidation state from +2 to +3 and back.

  Step 1 (Oxidation): \(2Fe^{2+} + S_2O_8^{2-} \rightarrow 2Fe^{3+} + 2SO_4^{2-}\)
  Step 2 (Reduction): \(2Fe^{3+} + 2I^- \rightarrow 2Fe^{2+} + I_2\)
  Overall: \(S_2O_8^{2-} + 2I^- \rightarrow 2SO_4^{2-} + I_2\) (The Fe²⁺ is regenerated.)

• Autocatalysis (Mn²⁺): This is a special case where a reaction is catalysed by one of its own products.
  In the titration of ethanedioate (\(C_2O_4^{2-}\)) with permanganate (\(MnO_4^-\)), the reaction is initially slow. However, the product, \(Mn^{2+}\), speeds up the reaction.
  Mechanism: \(Mn^{2+}\) is oxidised to \(Mn^{3+}\) by \(MnO_4^-\), and the resulting \(Mn^{3+}\) then reacts rapidly with \(C_2O_4^{2-}\) to reform \(Mn^{2+}\).