The Periodic Table - Transition elements (3)
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1.
Describe the transition elements as metals that: (c) form coloured compounds. Explain the electronic transitions that cause this colour, relating it to the d-orbitals.
Many transition metal compounds are coloured due to the presence of partially filled d-orbitals. When a transition metal compound absorbs light, electrons in the d-orbitals can be excited to higher energy levels. The wavelengths of light absorbed correspond to the energy differences between these d-orbital levels. Different transition metals and different oxidation states of the same metal have different energy differences between their d-orbitals, leading to the absorption of different wavelengths of light and therefore different colours. The colour we see is the complementary colour to the wavelengths absorbed. For example, a compound might absorb green light and appear blue.
d-orbital Energy Levels & Colour Absorption
| Transition Metal | Oxidation State | d-orbital Energy Level Difference | Colour Absorbed |
|---|
| Copper(II) | Cu2+ | Large | Green | | Iron(II) | Fe2+ | Small | Magenta |
|
2.
Describe transition elements as having ions with variable oxidation numbers. Include in your description the examples of iron(II) and iron(III), explaining how the variable oxidation states arise from the electronic configuration of the metal atom.
Transition elements are characterized by their ability to form ions with multiple oxidation states. This variability arises from the relatively small energy difference between the (n-1)d and ns orbitals. The electrons in these orbitals are readily lost to form cations with different positive charges. For example, iron (Fe) has an electronic configuration of [Ar] 3d6 4s2. The loss of electrons from the 4s orbital and/or the 3d orbitals results in different oxidation states.
Iron(II) (Fe2+) is formed when two electrons are removed from the 3d orbitals and one from the 4s orbital. This leaves a half-filled 3d subshell, which is a relatively stable configuration. The electronic configuration of Fe2+ is [Ar] 3d6.
Iron(III) (Fe3+) is formed when two electrons are removed from the 4s orbital and three electrons are removed from the 3d orbitals. This leaves a completely empty 3d subshell. The electronic configuration of Fe3+ is [Ar] 3d5. The ease with which electrons can be removed from the 4s and 3d orbitals explains the variable oxidation states of iron. The stability of the resulting ion is also a factor, with half-filled and completely filled subshells being particularly stable.
3.
Explain why transition metal compounds often have vibrant colours. Give examples of two transition metal compounds and state their colours.
Transition metal compounds often exhibit vibrant colours due to the presence of d-d electronic transitions. In a transition metal complex, the metal ion has unpaired electrons in its d orbitals. These unpaired electrons can absorb light of specific wavelengths when an electron transitions between different d orbitals. The wavelengths of light absorbed correspond to the colour we observe. The colour of a compound depends on the energy difference between the d orbitals and the nature of the ligands surrounding the metal ion. Different ligands cause different energy differences, leading to different colours.
Examples:
- [Fe(CN)6]3-: This complex is deep purple in colour. The d-d transition involves the movement of electrons between the d orbitals, which absorb light in the yellow-green region of the visible spectrum, resulting in the absorption of the complementary colour, purple.
- Cupric chloride (CuCl2): This compound is typically blue. The blue colour arises from the d-d transition involving the d orbitals of the Cu2+ ion.