# OCR Advanced GCE in Chemistry/Colour

First review the section on colour in the trends and patterns module:

## 3d Subshell

All five orbitals in the 3d subshell have two electrons each and are all at the same energy level (degenerate). However as we are about to find out this is only when the atom or ion involved is isolated (not bonded with anything else). Electrons will fill any of these orbitals in transitional atoms or ions and will fill each one before pairing.
d-orbital splitting: This is where the d-orbitals are in an (octahedral) transition metal complex so are no longer isolated. Six coordinate bonds are formed along the same axis as the 3d$_{z2}$ and 3d$_{x2y2}$ orbitals. Consequently, these orbitals are of higher energy than the other orbitals. This creates a higher tier energy level orbitals with an energy gap between them and the lower energy orbitals and is labelled ΔE.

## Electron promotion

If a transition element complex has at least one electron in a lower energy orbital and at least one space in a higher energy orbital the electron can jump to the higher energy orbital. This requires energy though, that's why light is needed. The energy gap ΔE previously described is the amount of energy from absorbed light required to allow a promotion. According to quantum theory and the equation e=hf, the energy of light is proportional to its wavelength. Therefore the amount of energy for a particular promotion will correspond to a certain wavelength (and therefore colour) of light. Transitional metals absorb wavelengths of light which are in humans visible range so we see colour.
If a particular wavelength of light is absorbed by the complex for electron promotion then the light reflected will be without that wavelength so will appear a different colour to white (all wavelengths).
Using the colour wheel from the trends and patterns module we can work out the colour of a complex solution that requires 4*10$^{-19}$J for a promotion:
green light has energy 4*10$^{-19}$J so green is absorbed and reflected light appears violet.
Compounds that do not fit the criteria explained above will be white e.g. Cu+ compounds are colourless because the electron configuration is [Ar]3d$^{10}$4s$^0$. As the 3d subshell is full there cannot be electron promotion.

## Changes in colour from ligand exchange

If ligands are substituted the E gap may increase or decrease, this will change the frequency of light absorbed from white light:
* If the energy gap becomes smaller the energy, and therefore frequency, of light absorbed will go down. This light will have a greater wavelength e.g. blue-->green light absorbed so yellow--->red viewed.
* If the energy gap increases the frequency of light needed increases. The wavelength of absorbed light decreases e.g. orange ---> blue absorbed, blue ---> yellow viewed
Remember: With all waves Velocity = Frequency X wavelength
Also, when a reactant is added to a complex a reduction or oxidation reaction could also occur with the transitional metal. This could also change the colour of the solution.

## Metal complex uses and spectra

Transitional elements are used for their colours in pigments for their brilliance in colour and for pure white:
* TiO$_2$ is used in white paint as it is excellent at hiding other colours beneath
* Phthalocyanine Blue BN is used for brilliant blue in paint and dyes - it is very stable
Visible spectrometry can be used to predict the colour of a transition metal complex. A graph can be drawn using a visible spectrometer for a complex and the relative absorbance of each wavelength can be identified.

## Questions

1. What can we call the orbitals of an isolated transition metal atom/ion?
2. Why does d-orbital splitting happen?
3. Which d-orbitals are the higher energy ones? draw them
4. What does the energy gap correspond to?
5. How can i change the energy gap and what will it do?
Answers