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!
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, Zn²⁺, has the configuration 3d¹⁰ (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, Sc³⁺, has the configuration 3d⁰ (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:
- When forming ions, TEs lose the 4s electrons first, even though they are filled before 3d. (e.g., Fe: [Ar] 3d⁶ 4s² $\rightarrow$ Fe²⁺: [Ar] 3d⁶).
- 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 3d$_{xy}$ orbital looks like a four-leaf clover lying flat on the x-y plane.
The 3d$_{z^2}$ orbital looks like a dumbbell shape along the z-axis with a "donut" (a ring of electron density) surrounding the middle.
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: V²⁺, V³⁺, VO²⁺ (V(IV)), and VO₃⁻ (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 (N₂ + 3H₂ $\rightleftharpoons$ 2NH₃).
- 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 ligands (the 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: Fe²⁺/Fe³⁺ in the reaction between I⁻ and S₂O₈²⁻.
3. Formation of Coloured Compounds
Transition metal ions in solution or in complexes are almost always vividly coloured.
- Explanation: This property is explained in detail in Section 28.3, but the short answer is that 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 don't form as many complexes).
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: H₂O (water), NH₃ (ammonia), Cl⁻ (chloride), CN⁻ (cyanide).
- Bidentate: Forms two dative bonds (two "teeth") to the same central ion.
- Examples: 1,2-diaminoethane (en, H₂NCH₂CH₂NH₂), ethanedioate ion (C₂O₄²⁻).
- Polydentate: Forms many dative bonds.
- Example: EDTA⁴⁻ (ethylenediaminetetraacetate ion) typically forms six 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: [Ag(NH₃)₂]⁺
- CN=4: Two possibilities:
- Tetrahedral (109.5°). Often seen with Cl⁻ ligands. Example: [CuCl₄]²⁻
- Square Planar (90°). Commonly seen with Pt(II) and Ni(II). Example: [Pt(NH₃)₂Cl₂]
- CN=6: Octahedral (90°). Most common shape. Example: [Cu(H₂O)₆]²⁺ or [Co(NH₃)₆]³⁺
Ligand Exchange Reactions
Ligands can be swapped, and this exchange often results in a dramatic colour change, which is easy to observe in the lab. These exchanges usually happen 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: [Cu(H₂O)₆]²⁺
- Adding Ammonia (NH₃):
A small amount of NH₃(aq) causes precipitation (a dirty blue solid, Cu(OH)₂).
$\text{[Cu(H₂O)₆]}^{2+}(\text{aq}) + 2\text{OH}^{-}(\text{aq}) \rightarrow \text{Cu(OH)}_{2}(\text{s}) + 6\text{H₂O}(\text{l})$
Adding excess NH₃ redissolves the precipitate to form a deep blue solution:
$\text{Cu(OH)}_{2}(\text{s}) + 4\text{NH}_{3}(\text{aq}) \rightarrow \text{[Cu(NH}_{3})_{4}\text{(H₂O)}_{2}]^{2+}(\text{aq}) + 2\text{OH}^{-}(\text{aq})$
- Adding Chloride Ions (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 (or sometimes square planar).
$\text{[Cu(H₂O)₆]}^{2+}(\text{aq}) + 4\text{Cl}^{-}(\text{aq}) \rightleftharpoons \text{[CuCl}_{4}]^{2-}(\text{aq}) + 6\text{H₂O}(\text{l})$
Colour change: Pale blue (octahedral) $\rightarrow$ yellow-green (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°) (See Topic 24.2).
Rule Reminder: A reaction is feasible if the E° of the reduction half-reaction is more positive than the E° of the oxidation half-reaction. Overall E° must be positive.
Key Reactions (Must Know!):
- Manganate(VII) / Oxalate (C₂O₄²⁻) in acid:
MnO₄⁻ is a powerful oxidising agent (Mn is reduced from +7 to +2). C₂O₄²⁻ is oxidised to CO₂. This reaction is autocatalysed by the product Mn²⁺.
$\text{2MnO}_{4}^{-}(\text{aq}) + 5\text{C}_{2}\text{O}_{4}^{2-}(\text{aq}) + 16\text{H}^{+}(\text{aq}) \rightarrow 2\text{Mn}^{2+}(\text{aq}) + 10\text{CO}_{2}(\text{g}) + 8\text{H}_{2}\text{O}(\text{l})$
- Manganate(VII) / Iron(II) (Fe²⁺) in acid:
MnO₄⁻ oxidises Fe²⁺ (Fe is oxidised from +2 to +3).
$\text{MnO}_{4}^{-}(\text{aq}) + 5\text{Fe}^{2+}(\text{aq}) + 8\text{H}^{+}(\text{aq}) \rightarrow \text{Mn}^{2+}(\text{aq}) + 5\text{Fe}^{3+}(\text{aq}) + 4\text{H}_{2}\text{O}(\text{l})$
- Copper(II) / Iodide (I⁻):
Copper(II) ions oxidise iodide ions to iodine (I is oxidised from -1 to 0). Copper(II) is reduced from +2 to +1 (precipitated as solid CuI).
$\text{2Cu}^{2+}(\text{aq}) + 4\text{I}^{-}(\text{aq}) \rightarrow 2\text{CuI}(\text{s}) + \text{I}_{2}(\text{aq})$
28.3 Colour of Complexes
The colour arises from the way the ligands interact with the d orbitals of the central metal ion. Don't worry if this seems complicated—we will break it down!
1. Degenerate and Splitting of d Orbitals
In an isolated transition metal atom or ion, the five d orbitals ($d_{xy}, d_{yz}, d_{xz}, d_{x^2-y^2}, d_{z^2}$) all have the same energy. They are called degenerate.
When ligands approach the central ion (forming a complex), the d orbitals are no longer degenerate. The repulsion between the electron lone pairs on the ligands and the electrons in the d orbitals causes the d orbitals to split into two groups with different energies. 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
When white light (which contains all visible frequencies) hits the complex ion, an electron from a lower energy d orbital absorbs a specific frequency of light energy equal to $\Delta E$ and is promoted to the higher energy d orbital. This is called a d-d transition.
The frequency of light absorbed is removed from the white light spectrum. The remaining light is transmitted or reflected, and this is the complementary colour that we observe.
- Example: If a complex absorbs yellow light, we see the complementary colour, blue.
3. Factors Affecting Colour
The magnitude of $\Delta E$ dictates which frequency is absorbed, and this depends on the ligand.
- Ligands: Different ligands cause different amounts of splitting ($\Delta E$). Strong field ligands (like CN⁻ or NH₃) cause a large $\Delta E$ and absorb high-energy, high-frequency light (blue/UV). Weak field ligands (like H₂O or Cl⁻) cause a small $\Delta E$ and absorb low-energy, low-frequency light (red/IR).
- Oxidation State: Higher oxidation states generally lead to smaller ions, resulting in stronger electrostatic attraction and greater $\Delta E$.
- Geometry: The shape (octahedral vs. tetrahedral) significantly changes the splitting pattern and therefore the $\Delta E$.
Ligand Exchange and Colour (Revisited):
When you swap H₂O ligands for NH₃ ligands in a copper complex:
$\text{[Cu(H₂O)₆]}^{2+}(\text{aq}) \quad \rightarrow \quad \text{[Cu(NH}_{3})_{4}\text{(H₂O)}_{2}]^{2+}(\text{aq})$
Colour change: Pale blue $\rightarrow$ Deep blue. NH₃ is a stronger ligand than H₂O, so it causes a larger $\Delta E$, meaning a different frequency (and thus a different complementary colour) is observed.
28.4 Stereoisomerism in Complexes
Complexes can exhibit stereoisomerism, specifically geometrical (cis/trans) and optical isomerism, particularly when they involve bidentate ligands.
Geometrical Isomerism (cis/trans)
This occurs when ligands occupy different positions relative to one another in space. It happens primarily in square planar and octahedral complexes.
- Square Planar Complexes (e.g., [Pt(NH₃)₂Cl₂]):
- cis-isomer: Identical ligands are adjacent (90° apart). Cisplatin is a famous example (anti-cancer drug).
- trans-isomer: Identical ligands are opposite (180° apart).
- Octahedral Complexes (e.g., [Co(NH₃)₄(H₂O)₂]²⁺):
- cis-isomer: The two unique ligands (H₂O in this case) are adjacent (90° apart).
- trans-isomer: The two unique ligands are opposite (180° apart).
Did you know? The difference between cis- and trans-isomers can be biologically vital. Cisplatin works as a chemotherapy agent, but the trans-isomer is inactive and toxic.
Optical Isomerism
This occurs in complexes that are non-superimposable mirror images of each other (like hands). This usually requires bidentate ligands in an octahedral arrangement.
- Example: $\text{[Ni}(\text{H}_{2}\text{NCH}_{2}\text{CH}_{2}\text{NH}_{2}\text{)}_{3}]^{2+}$ or $\text{[Ni(en)}_{3}]^{2+}$ (where 'en' is the bidentate ligand 1,2-diaminoethane). The three 'en' ligands wrap around the central nickel ion, creating two non-superimposable mirror images (enantiomers).
Overall Polarity
The overall polarity of a complex depends on its shape and the arrangement of its ligands.
- cis-isomers usually have a net dipole moment and are polar (due to asymmetric arrangement).
- trans-isomers often have their bond dipoles cancel out symmetrically and are usually non-polar.
28.5 Stability Constants, K$_{stab}$
When a complex ion forms in a solvent, it's an equilibrium process. The stability constant, K$_{stab}$, tells us how strong the bond is between the metal ion and the ligands.
Definition: The stability constant, $K_{stab}$, is the equilibrium constant for the formation of a complex ion in a solvent from its constituent ions or molecules.
Writing K$_{stab}$ Expressions
Consider the formation of the complex $\text{[Cu(NH}_{3})_{4}\text{]}^{2+}$ from $\text{Cu}^{2+}$ and $\text{NH}_{3}$ (in water):
$\text{Cu}^{2+}(\text{aq}) + 4\text{NH}_{3}(\text{aq}) \rightleftharpoons \text{[Cu(NH}_{3})_{4}\text{]}^{2+}(\text{aq})$
The stability constant expression is:
$$K_{stab} = \frac{[ \text{[Cu(NH}_{3})_{4}\text{]}^{2+}]}{[\text{Cu}^{2+}] [\text{NH}_{3}]^4}$$
Crucial Point: We exclude the concentration of water ($\text{[H}_{2}\text{O}]$) from the $K_{stab}$ expression, as it is considered a constant high concentration (like in $K_c$ expressions for aqueous solutions).
Using K$_{stab}$ to Explain Ligand Exchange
A large value of $K_{stab}$ indicates that the position of equilibrium lies far to the right, meaning the complex ion is very stable (favoured).
Ligand exchange reactions occur when the new ligand forms a complex with a larger $K_{stab}$ than the original complex. The more stable complex effectively displaces the less stable one.
- Example: The reaction of [Cu(H₂O)₆]²⁺ with NH₃ to form [Cu(NH₃)₄(H₂O)₂]²⁺ proceeds because the $\text{Cu}^{2+}/\text{NH}_{3}$ complex has a much larger $K_{stab}$ than the original $\text{Cu}^{2+}/\text{H}_{2}\text{O}$ complex. Ammonia forms a more stable complex, driving the reaction forward.
✅ Key Takeaways for Transition Elements
- Definition: TEs are d-block elements forming stable ions with incomplete d orbitals (3d¹ to 3d⁹).
- Properties Source: Similar energies of 3d and 4s sub-shells allow for variable O.S. and accessibility of vacant d orbitals.
- Ligands: Donate lone pairs to the metal ion via dative covalent bonds to form complexes.
- Colour: Caused by the d-d transition when electrons absorb light energy ($\Delta E$) corresponding to the splitting of d orbitals.
- Stability: A complex with a large $K_{stab}$ is very stable, which is why it displaces weaker ligands in exchange reactions.
- Manganate(VII) / Oxalate (C₂O₄²⁻) in acid: