Chemistry of transition elements (A Level only)

Chemistry of Transition Elements (Topic 28)

Hey there! Welcome to one of the most colourful and interesting topics in A-Level Chemistry. Transition elements (TEs) are the weird and wonderful elements in the middle of the Periodic Table that give us bright paints, powerful magnets, and essential biological catalysts. This chapter ties together concepts from AS chemistry (like redox and orbitals) with brand-new ideas about complex structures and light absorption. Let's dive in!

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28.1 General Physical and Chemical Properties

Defining a Transition Element

The term "d-block element" refers to elements in groups 3 to 12. However, the definition of a true transition element is more specific:

Definition: A transition element is a d-block element which forms one or more stable ions with incomplete d orbitals (i.e., partially filled d sub-shells).

• The first row of transition elements studied runs from Titanium (Ti) to Copper (Cu).
• Zinc (Zn) is usually excluded because its ion, \(\text{Zn}^{2+}\), has the configuration \([\text{Ar}]\,3\text{d}^{10}\) (a full d sub-shell). Since it doesn't have an incomplete d orbital in its stable ion, it's a d-block element, but not a transition element.
• Scandium (Sc) is also excluded because its only stable ion, \(\text{Sc}^{3+}\), has the configuration \([\text{Ar}]\,3\text{d}^0\) (an empty d sub-shell).

Electron Configuration: The Key to Weirdness

The special properties of TEs stem from the fact that the 3d and 4s sub-shells have very similar energies. This causes two important effects:

1. When forming ions, TEs lose the 4s electrons first, even though they are filled before 3d. (e.g., \(\text{Fe}: [\text{Ar}]\,3\text{d}^6\,4\text{s}^2 \rightarrow \text{Fe}^{2+}: [\text{Ar}]\,3\text{d}^6\)).
2. Since 3d and 4s are close in energy, it takes very little energy to remove or share varying numbers of electrons, leading to variable oxidation states.

Quick Review: Orbital Shapes
You need to be able to recall the shapes of certain d orbitals:
The \(3\text{d}_{xy}\) orbital looks like a four-leaf clover lying flat on the x-y plane.
The \(3\text{d}_{z^2}\) orbital looks like a dumbbell shape along the z-axis with a "donut" (a ring of electron density) surrounding the middle.

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28.1 The Four Characteristic Properties

Transition elements are famous for four linked characteristics:

1. Variable Oxidation States (O.S.)

Transition elements can exist in many different oxidation states (e.g., Mn ranges from +2 to +7).

• Explanation: Due to the close energy levels of the 3d and 4s sub-shells, different numbers of electrons can be readily lost or shared.
• Example: Vanadium forms stable ions: \(\text{V}^{2+}\), \(\text{V}^{3+}\), \(\text{VO}^{2+}\) (V(IV)), and \(\text{VO}_3^-\)/\(\text{VO}_2^+\) (V(V)).

2. Behaviour as Catalysts

Many TEs and their compounds are excellent catalysts, crucial in industrial processes.

• Heterogeneous Catalysis: The TE is in a different phase from the reactants (usually solid metal/oxide, reactants are gas/liquid).
  • Example: Iron (Fe) in the Haber process (\(\text{N}_2 + 3\text{H}_2 \rightleftharpoons 2\text{NH}_3\)).
  • Mode of Action: The reactant molecules adsorb (stick) onto the surface of the metal, weakening their bonds. The metal provides vacant d orbitals to form dative bonds with reactants. The product then desorbs.
• Homogeneous Catalysis: The TE and reactants are in the same phase (usually aqueous solution).
  • Explanation: The TE can change its oxidation state easily, allowing it to provide an alternative reaction pathway with lower activation energy (\(E_a\)).
  • Example: \(\text{Fe}^{2+}/\text{Fe}^{3+}\) in the reaction between \(\text{I}^-\) and \(\text{S}_2\text{O}_8^{2-}\).

3. Formation of Coloured Compounds

Transition metal ions in solution or in complexes are almost always vividly coloured.

• Explanation: The partial filling of the d-orbitals allows electrons to absorb specific frequencies of visible light and jump to higher energy d-orbitals (known as d-d transitions). The colour we see is the light that is not absorbed (the complementary colour).

4. Formation of Complex Ions

Transition metal ions readily form complex ions when bonded to molecules or ions called ligands.

• Explanation: TE ions are small, highly charged cations with vacant d orbitals (that are energetically accessible), which can readily accept lone pairs of electrons from ligands to form dative covalent bonds (coordinate bonds).

Common Mistake Alert!
The reason TEs form complex ions is the availability of vacant d orbitals, which are used to accept electron pairs from ligands. It is NOT simply because they are small and highly charged (Group 13 ions are also small and charged but do not form complexes in the same way).

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28.2 Complex Ions, Ligands, and Redox Chemistry

Ligands and Complexes

A ligand is a species (a molecule or an ion) that possesses at least one lone pair of electrons that can be donated to a central metal atom/ion to form a dative covalent bond.

A complex (or complex ion) is a molecule or ion formed by a central metal atom or ion surrounded by one or more ligands.

The number of dative bonds formed to the central metal ion is the coordination number (CN).

Types of Ligands

• Monodentate: Forms one dative bond (one "tooth").
  • Examples: \(\text{H}_2\text{O}\) (water), \(\text{NH}_3\) (ammonia), \(\text{Cl}^-\) (chloride), \(\text{CN}^-\) (cyanide).
• Bidentate: Forms two dative bonds (two "teeth") to the same central ion.
  • Examples: 1,2-diaminoethane (en, \(\text{H}_2\text{NCH}_2\text{CH}_2\text{NH}_2\)), ethanedioate ion (\(\text{C}_2\text{O}_4^{2-}\)).
• Polydentate: Forms many dative bonds.
  • Example: \(\text{EDTA}^{4-}\) (ethylenediaminetetraacetate ion) forms six dative bonds (CN=6).

Analogy: Think of the metal ion as your head, and ligands as hands. Monodentate ligands have one hand to grab your head; bidentate ligands have two hands (a stronger grip!).

Geometry of Complexes (Shapes)

The shape of the complex depends on the Coordination Number (CN).

• CN=2: Linear (180°). Example: \([\text{Ag}(\text{NH}_3)_2]^+\)
• CN=4: Two possibilities:
  • Tetrahedral (109.5°). Often seen with large ligands like \(\text{Cl}^-\). Example: \([\text{CuCl}_4]^{2-}\), \([\text{CoCl}_4]^{2-}\)
  • Square Planar (90°). Commonly seen with Pt(II) and Ni(II). Example: \([\text{Pt}(\text{NH}_3)_2\text{Cl}_2]\)
• CN=6: Octahedral (90°). Most common shape. Example: \([\text{Cu}(\text{H}_2\text{O})_6]^{2+}\) or \([\text{Co}(\text{NH}_3)_6]^{3+}\)

Ligand Exchange Reactions

Ligands can be swapped, and this exchange often results in a dramatic colour change. These exchanges usually occur when the new ligand forms a more stable complex (see 28.5).

Example 1: Copper(II) (CN=6, Octahedral)

Start: Aqueous copper(II) ions are pale blue: \([\text{Cu}(\text{H}_2\text{O})_6]^{2+}\)

1. Adding Ammonia (\(\text{NH}_3\)):

   A small amount of \(\text{NH}_3(\text{aq})\) causes precipitation of a pale blue solid, \(\text{Cu}(\text{OH})_2\):

   \([\text{Cu}(\text{H}_2\text{O})_6]^{2+}(\text{aq}) + 2\text{OH}^-(\text{aq}) \rightarrow \text{Cu}(\text{OH})_2(\text{s}) + 6\text{H}_2\text{O}(\text{l})\)

   Adding excess \(\text{NH}_3\) redissolves the precipitate to form a deep blue solution:

   \([\text{Cu}(\text{H}_2\text{O})_6]^{2+}(\text{aq}) + 4\text{NH}_3(\text{aq}) \rightarrow [\text{Cu}(\text{NH}_3)_4(\text{H}_2\text{O})_2]^{2+}(\text{aq}) + 4\text{H}_2\text{O}(\text{l})\)

2. Adding Chloride Ions (\(\text{Cl}^-\)) / Concentrated HCl:

   Chloride ions are larger than water molecules, so the coordination number reduces from 6 to 4, and the shape changes from octahedral to tetrahedral.

   \([\text{Cu}(\text{H}_2\text{O})_6]^{2+}(\text{aq}) + 4\text{Cl}^-(\text{aq}) \rightleftharpoons [\text{CuCl}_4]^{2-}(\text{aq}) + 6\text{H}_2\text{O}(\text{l})\)

   Colour change: Pale blue (octahedral) \(\rightarrow\) yellow-green (tetrahedral).

Example 2: Cobalt(II) (CN=6, Octahedral)

Start: Aqueous cobalt(II) ions are pink: \([\text{Co}(\text{H}_2\text{O})_6]^{2+}\)

1. Adding Ammonia (\(\text{NH}_3\)):

   Adding excess \(\text{NH}_3(\text{aq})\) replaces water ligands with ammonia ligands:

   \([\text{Co}(\text{H}_2\text{O})_6]^{2+}(\text{aq}) + 6\text{NH}_3(\text{aq}) \rightarrow [\text{Co}(\text{NH}_3)_6]^{2+}(\text{aq}) + 6\text{H}_2\text{O}(\text{l})\)

   Colour change: Pink \(\rightarrow\) straw / brown (which slowly oxidises in air to yellow-brown \([\text{Co}(\text{NH}_3)_6]^{3+}\)).

2. Adding Chloride Ions (\(\text{Cl}^-\)) / Concentrated HCl:

   \([\text{Co}(\text{H}_2\text{O})_6]^{2+}(\text{aq}) + 4\text{Cl}^-(\text{aq}) \rightleftharpoons [\text{CoCl}_4]^{2-}(\text{aq}) + 6\text{H}_2\text{O}(\text{l})\)

   Colour change: Pink (octahedral) \(\rightarrow\) blue (tetrahedral).

Redox Reactions Involving Transition Elements

Since TEs have variable oxidation states, they are excellent oxidising or reducing agents. You need to be able to predict the feasibility of these reactions using standard electrode potentials (\(E^\circ\)).

Key Redox Reactions:

1. Manganate(VII) / Ethanedioate (\(\text{C}_2\text{O}_4^{2-}\)) in acid:

   \(\text{MnO}_4^-\) oxidises \(\text{C}_2\text{O}_4^{2-}\) to \(\text{CO}_2\), and Mn is reduced from +7 to +2. This reaction is autocatalysed by the product \(\text{Mn}^{2+}\).

   \(2\text{MnO}_4^-(\text{aq}) + 5\text{C}_2\text{O}_4^{2-}(\text{aq}) + 16\text{H}^+(\text{aq}) \rightarrow 2\text{Mn}^{2+}(\text{aq}) + 10\text{CO}_2(\text{g}) + 8\text{H}_2\text{O}(\text{l})\)

2. Manganate(VII) / Iron(II) (\(\text{Fe}^{2+}\)) in acid:

   \(\text{MnO}_4^-\) oxidises \(\text{Fe}^{2+}\) to \(\text{Fe}^{3+}\).

   \(\text{MnO}_4^-(\text{aq}) + 5\text{Fe}^{2+}(\text{aq}) + 8\text{H}^+(\text{aq}) \rightarrow \text{Mn}^{2+}(\text{aq}) + 5\text{Fe}^{3+}(\text{aq}) + 4\text{H}_2\text{O}(\text{l})\)

3. Copper(II) / Iodide (\(\text{I}^-\)):

   Copper(II) ions oxidise iodide ions to iodine (\(\text{I}_2\)), while copper(II) is reduced to copper(I), forming an off-white precipitate of \(\text{CuI}\).

   \(2\text{Cu}^{2+}(\text{aq}) + 4\text{I}^-(\text{aq}) \rightarrow 2\text{CuI}(\text{s}) + \text{I}_2(\text{aq})\)

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28.3 Colour of Complexes

1. Degenerate and Splitting of d Orbitals

In an isolated transition metal atom or ion, the five d orbitals (\(3\text{d}_{xy}, 3\text{d}_{yz}, 3\text{d}_{xz}, 3\text{d}_{x^2-y^2}, 3\text{d}_{z^2}\)) all have the same energy. They are called degenerate.

When ligands approach the central ion (forming a complex), electrostatic repulsion between the ligand lone pairs and the d electrons causes the d orbitals to split into two groups with different energy levels. This energy difference is called the crystal field splitting energy (\(\Delta E\)).

The splitting pattern depends on the complex geometry:

• Octahedral (CN=6): Splits into two higher energy d orbitals and three lower energy d orbitals. (2 high, 3 low)
• Tetrahedral (CN=4): Splits into three higher energy d orbitals and two lower energy d orbitals. (3 high, 2 low)

2. The Origin of Colour & the \(\Delta E\) Relationship

When white light hits the complex ion, an electron from a lower energy d orbital absorbs a photon of light energy equal to \(\Delta E\) and is promoted to a higher energy d orbital. This is called a d-d transition.

The relationship between the energy absorbed \(\Delta E\), frequency \(\nu\), and wavelength \(\lambda\) is given by:

\(\Delta E = h\nu = \frac{hc}{\lambda}\)

where \(h\) is Planck's constant and \(c\) is the speed of light.

The absorbed wavelength is removed from the visible spectrum. The remaining transmitted or reflected wavelengths give the complementary colour observed (e.g., absorbing red light results in a blue/cyan appearance).

3. Factors Affecting Colour

• Nature of Ligands: Different ligands cause different amounts of splitting (\(\Delta E\)), changing the wavelength \(\lambda\) absorbed.
• Oxidation State: Higher oxidation states increase \(\Delta E\).
• Geometry: Splitting differs between octahedral and tetrahedral shapes.

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28.4 Stereoisomerism in Complexes

Complexes can exhibit stereoisomerism, specifically geometrical (cis/trans) and optical isomerism.

Geometrical Isomerism (cis/trans)

• Square Planar Complexes (e.g., \([\text{Pt}(\text{NH}_3)_2\text{Cl}_2]\)):
  • cis-isomer: Identical ligands are adjacent (90° apart). Cisplatin is a vital anti-cancer drug.
  • trans-isomer: Identical ligands are opposite (180° apart).
• Octahedral Complexes (e.g., \([\text{Co}(\text{NH}_3)_4(\text{H}_2\text{O})_2]^{2+}\)):
  • cis-isomer: The two \(\text{H}_2\text{O}\) ligands are adjacent (90° apart).
  • trans-isomer: The two \(\text{H}_2\text{O}\) ligands are opposite (180° apart).

Mechanism of Action of Cisplatin:
Cisplatin acts as an anti-cancer drug by entering cells and binding to DNA. The chlorine ligands are replaced, allowing the platinum atom to bind datively to nitrogen atoms on guanine bases in DNA. This cross-links the DNA strands, distorting the DNA helix, which prevents DNA replication and cell division, ultimately triggering cell death.

Optical Isomerism

Optical isomers are non-superimposable mirror images of each other (enantiomers). This commonly occurs in octahedral complexes containing bidentate ligands.

• Example: \([\text{Ni}(\text{en})_3]^{2+}\) (where 'en' is 1,2-diaminoethane). The three bidentate ligands wrap around the central nickel ion to produce two non-superimposable mirror images.

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28.5 Stability Constants, \(K_{\text{stab}}\)

Definition: The stability constant, \(K_{\text{stab}}\), is the equilibrium constant for the formation of a complex ion in a solvent from its constituent ions or molecules.

Writing \(K_{\text{stab}}\) Expressions

Consider the formation of \([\text{Cu}(\text{NH}_3)_4]^{2+}\) from \(\text{Cu}^{2+}\) and \(\text{NH}_3\) in aqueous solution:

\(\text{Cu}^{2+}(\text{aq}) + 4\text{NH}_3(\text{aq}) \rightleftharpoons [\text{Cu}(\text{NH}_3)_4]^{2+}(\text{aq})\)

The stability constant expression is:

\(K_{\text{stab}} = \frac{[[\text{Cu}(\text{NH}_3)_4]^{2+}]}{[\text{Cu}^{2+}][\text{NH}_3]^4}\)

Crucial Point: We exclude the concentration of water (\([\text{H}_2\text{O}]\)) from the \(K_{\text{stab}}\) expression because water is the solvent and its concentration remains essentially constant.

Using \(K_{\text{stab}}\) to Explain Ligand Exchange

A large value of \(K_{\text{stab}}\) indicates that the position of equilibrium lies far to the right, meaning the complex ion is very stable.

Ligand exchange reactions occur when the incoming ligand forms a complex with a larger \(K_{\text{stab}}\) than the original complex. The more stable complex displaces the less stable one.