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A-Level Properties and complexes of first-row transition metals
What the A-Level syllabus expects for Properties and complexes of first-row transition metals, and how to practise it.
What the syllabus expects
- Set out the definition of a transition element, a d-block species whose atoms, or whose cations, keep a d subshell that stays only partly filled.
- For a first-row transition element, write the electron arrangement of the atom and of each ion it forms.
Scope: see also 1(h) - Explain why, across the transition elements, both the first ionisation energy and the atomic radius stay very nearly constant.
- Compare, in qualitative terms, the melting points and densities of the transition elements against calcium as a representative s-block metal.
- Describe how readily transition elements adopt more than one oxidation state.
- From a supplied electron configuration, predict which oxidation states a transition element is likely to show.
- Describe and account for the redox couples Fe3+/Fe2+, then MnO4-/Mn2+, and also Cr2O72-/Cr3+, as worked examples.
Scope: see also Section 12 - Use E-standard values to judge how likely a redox reaction is.
Scope: see also 12(f)(ii) - Define ligand and complex, illustrated by copper(II) ions bound to water, ammonia and chloride ligands.
Scope: this takes in the transition-metal complexes listed in the Qualitative Analysis Notes - Explain qualitatively how ligands can swap over, as when the complexes of (i) form with their attendant colour changes, and as in the CO/O2 exchange within haemoglobin.
- Draw on how the d orbitals are shaped and oriented to describe how, in an octahedral complex, a once-degenerate d-orbital set splits into two energy levels.
- Explain why transition-element complexes are usually coloured by appealing to d-orbital splitting and to d-d transitions.
Scope: no need to know how ligand-field strength is ranked - Explain how certain transition elements or their compounds can act as catalysts.
Scope: see also 8(j)
How it's examined
Questions on this topic most often ask you to explain, name, sketch. About 8% of the past-paper style questions in Rae's bank for this subject sit in this topic.
Worked examples
Example 1 (5 marks)
(c)(ii) [Co(NH3)6]3+ is yellow-brown in colour, whereas [CoF6]3- is blue. Explain the reason for this difference in colour. [5] [Total: 20]
Show the worked answer
In an octahedral complex the ligands split the five 3d orbitals into a lower set (t2g) and a higher set (eg) separated by an energy gap ΔE (the ligand-field splitting). Colour arises because an electron absorbs a photon of visible light and is promoted from the lower to the higher d level (a d-d transition); the complementary colour of the light absorbed is transmitted and seen. The size of ΔE depends on the field strength of the ligand: from the spectrochemical series NH3 is a stronger-field ligand than F-, so ΔE is larger for [Co(NH3)6]3+ than for [CoF6]3-. Since ΔE = hc/λ, the larger gap means [Co(NH3)6]3+ absorbs shorter-wavelength (higher-energy) light while [CoF6]3- absorbs longer-wavelength light. Absorbing different wavelengths means different complementary colours are transmitted, so one appears yellow-brown and the other blue.
Example 2 (3 marks)
Name a transition-metal cation that could be used to tell apart the oxidising powers of Br₂ and I₂. Support your choice with suitable reasoning and calculations.
Show the worked answer
Use the Fe3+/Fe2+ couple, E(Fe3+/Fe2+) = +0.77 V, which lies between E(Br2/Br-) = +1.07 V and E(I2/I-) = +0.54 V. Consider whether each halogen can oxidise Fe2+ to Fe3+. With Br2: Ecell = E(Br2/Br-) - E(Fe3+/Fe2+) = 1.07 - 0.77 = +0.30 V (>0), so Br2 does oxidise Fe2+ to Fe3+ (feasible). With I2: Ecell = E(I2/I-) - E(Fe3+/Fe2+) = 0.54 - 0.77 = -0.23 V (<0), so I2 does not oxidise Fe2+ (not feasible). Because Br2 oxidises the iron(II) but iodine does not, the Fe3+/Fe2+ system distinguishes the two halogens' oxidising powers (Br2 is the stronger oxidant).
Example 3 (2 marks)
Explain why TiCl3 gives a violet solution whereas TiCl4 gives a colourless one.
Show the worked answer
Colour in transition-metal species arises from d-d electronic transitions, which require partially filled d-orbitals. In TiCl3 the metal is Ti3+, which is [Ar]3d1: it has one d electron that can be promoted between the split d-orbital levels, absorbing part of the visible spectrum and giving the complementary colour (violet). In TiCl4 the metal is Ti4+, which is [Ar]3d0: with no d electrons, no d-d transition is possible, so no visible light is absorbed and the solution is colourless.
More worked questions on this topic
- When ligands coordinate to a copper(II) ion and cause the d-orbitals to split in energy, the re (2 marks)
- Under suitable conditions, [V(H2O)6]3+ reacts with zinc to give a solution of a different colou (2 marks)
- In a free Fe2+ ion the five 3d orbitals are degenerate. Sketch an energy-level diagram contrast (2 marks)
- Table 4.2 records some physical data for calcium, cobalt and manganese (relative atomic masses (2 marks)
More A-Level H2 Chemistry topics
The make-up of the atom and how its electrons are arranged · How atoms bond and how that governs a substance's behaviour · Ideal gases and working with gas mixtures · Competing definitions of acids and bases · Trends in the elements across a period and down a group · The mole and reacting-quantity calculations · all of A-Level H2 Chemistry