Revise General characteristic chemical properties of the first set of transition elements, titanium to copper for Cambridge 9701 (28.2)

A defining characteristic of transition elements is the formation of coloured compounds and complex ions in solution. This phenomenon arises from the presence of a partially filled d-subshell. When ligands bond to the central metal ion, they cause the five degenerate d-orbitals to split into two sets of different energy levels. An electron can be promoted from a lower energy d-orbital to a higher energy d-orbital by absorbing energy from the visible light spectrum. The energy absorbed corresponds to a specific frequency (ΔE = hf). The colour we perceive is the complementary colour, composed of the frequencies of light that are transmitted or reflected. The specific colour depends on the energy gap (ΔE), which is influenced by the metal's identity, its oxidation state, and the nature of the ligands.

Transition elements from titanium to copper exhibit variable oxidation states in their compounds. This is because the 4s and 3d sub-shells are very close in energy. Consequently, a variable number of electrons can be lost or shared during chemical reactions with relatively small successive increases in ionisation energy. For example, manganese can exist in all oxidation states from +2 (e.g., Mn²⁺) to +7 (e.g., MnO₄⁻). The stability of these oxidation states varies; for iron, the +2 and +3 states are common, with Fe³⁺ being particularly stable due to its half-filled d-subshell (3d⁵ configuration). This ability to readily change oxidation state makes transition elements and their ions effective oxidising and reducing agents, underpinning their role in redox titrations and catalysis.

A ligand substitution reaction is one in which a ligand in a complex ion is replaced by another ligand. These reactions are often reversible equilibria. A classic example is the stepwise replacement of water ligands in the pale blue aqueous copper(II) ion, [Cu(H₂O)₆]²⁺. Adding concentrated hydrochloric acid replaces the water ligands with chloride ions, forming the yellow-green tetrahedral complex, [CuCl₄]²⁻. Alternatively, adding aqueous ammonia initially precipitates blue Cu(OH)₂, which then redissolves in excess ammonia to form the deep blue square planar complex, [Cu(NH₃)₄(H₂O)₂]²⁺. The position of the equilibrium is determined by the relative concentrations of the ligands and the relative stability of the complexes formed. Such reactions are almost always accompanied by a distinct colour change, making them easy to observe.

Transition elements and their compounds are highly effective catalysts in both industrial and biological systems. Their catalytic prowess stems from two key properties: their ability to exist in variable oxidation states and their ability to provide a surface onto which reactants can be adsorbed. In homogeneous catalysis, the catalyst is in the same phase as the reactants. For example, Fe²⁺(aq) catalyses the reaction between I⁻(aq) and S₂O₈²⁻(aq) by providing an alternative reaction pathway with a lower activation energy, cycling between the Fe²⁺ and Fe³⁺ oxidation states. In heterogeneous catalysis, the catalyst is in a different phase. For example, solid iron in the Haber process provides an active surface for the adsorption and bond weakening of N₂ and H₂ molecules.

A complex ion consists of a central metal ion surrounded by several ligands. A ligand is a molecule or ion with a lone pair of electrons that it can donate to the central metal ion, forming a coordinate (or dative covalent) bond. In this interaction, the metal ion acts as a Lewis acid (electron pair acceptor) and the ligand acts as a Lewis base (electron pair donor). The total number of coordinate bonds to the central ion is its coordination number. Common coordination numbers are 6 (forming an octahedral complex) and 4 (forming a tetrahedral or square planar complex).

The ligands in a complex are not permanently fixed; they can be replaced by other ligands in a ligand substitution reaction. These reactions are often reversible equilibria. The position of the equilibrium depends on factors like the relative stability of the complexes and the concentration of the ligands. A classic example is the reaction of the pale blue hexaquacopper(II) ion with chloride ions, which produces the yellow-green tetrachlorocuprate(II) ion. These reactions are visually striking and are a cornerstone of transition metal chemistry.

A hallmark of transition elements is their ability to exist in multiple, stable oxidation states. This arises because the 3d and 4s orbitals are very close in energy, allowing for the removal of a variable number of electrons. This property makes them central to redox chemistry. Two of the most important reagents in A-Level chemistry, potassium manganate(VII) ($KMnO_4$) and potassium dichromate(VI) ($K_2Cr_2O_7$), are powerful oxidising agents precisely because the transition metal is in a high oxidation state ($Mn^{VII}$, $Cr^{VI}$) and can readily be reduced.

Reduction of Manganate(VII) in acid: $MnO_4^-(aq) + 8H^+(aq) + 5e^- \rightarrow Mn^{2+}(aq) + 4H_2O(l)$

Reduction of Dichromate(VI) in acid: $Cr_2O_7^{2-}(aq) + 14H^+(aq) + 6e^- \rightarrow 2Cr^{3+}(aq) + 7H_2O(l)$

Transition elements and their compounds are exceptional catalysts. Their catalytic prowess stems from their ability to bring reactants together (adsorption in heterogeneous catalysis) and their capacity to cycle through different oxidation states (in homogeneous catalysis), providing a lower-energy reaction pathway. In heterogeneous catalysis, such as the use of solid iron in the Haber process, the d-orbitals on the metal surface form temporary bonds with reactant molecules, weakening their internal bonds and increasing the likelihood of a successful collision. In homogeneous catalysis, such as $Fe^{2+}(aq)$ catalysing the reaction between $S_2O_8^{2-}(aq)$ and $I^-(aq)$, the catalyst is oxidised in one step and then reduced back to its original state in a second step, completing the catalytic cycle.

Q: Why are zinc and scandium not considered typical transition elements chemically?

A: Zinc only forms the $Zn^{2+}$ ion, which has a full d-subshell ($3d^{10}$). Scandium only forms the $Sc^{3+}$ ion, which has an empty d-subshell ($3d^0$). Since neither forms stable ions with a partially filled d-subshell, they do not exhibit the characteristic properties like variable oxidation states or coloured compounds.

Q: Is there a difference between a coordinate bond and a dative covalent bond?

A: No, they are two different names for the same type of bond. A coordinate bond is a covalent bond where one atom provides both of the shared electrons. The term 'coordinate bond' is most commonly used in the context of transition metal complexes.

Q: How can I predict the colour of a complex ion?

A: You are not expected to predict the exact colour from first principles at A-Level. However, you must learn the specific colours of common complexes, such as the blue of $[Cu(H_2O)_6]^{2+}$, the yellow-green of $[CuCl_4]^{2-}$, the pink of $[Co(H_2O)_6]^{2+}$, and the green of $Cr^{3+}(aq)$. The colour depends on the metal ion, its oxidation state, and the ligands attached.