UV–Visible Spectroscopy

Why is a carrot orange but a potato white? Why does a solution of copper sulfate look blue, and why does it turn a much deeper blue when ammonia is added? Why can a biologist measure the amount of DNA in a tiny drop of liquid by shining ultraviolet light through it? All of these questions are answered by UV–visible spectroscopy, which studies how molecules absorb ultraviolet and visible light.

The region of the spectrum

UV–visible spectroscopy covers wavelengths from about 200 nm to 800 nm:

• Ultraviolet (UV): about 200–400 nm (below about 200 nm, oxygen in air absorbs strongly, so ordinary instruments don’t go lower)
• Visible: about 400–700 nm, from violet to red

Radiation in this range carries enough energy to move electrons from one energy level to a higher one. That’s why it’s called electronic spectroscopy.

What happens when light is absorbed

A molecule absorbs a photon when the photon’s energy exactly matches the gap between two electronic energy levels:

ΔE = hf = hc ÷ λ

An electron jumps from a lower-energy orbital (usually a bonding or non-bonding orbital) to a higher-energy, empty orbital (usually an antibonding orbital). The main types of transition:

Transition	Where it happens	Typical wavelength
σ → σ*	single bonds (C–C, C–H)	below 200 nm (not seen in ordinary UV–vis)
n → σ*	atoms with lone pairs (O, N, Cl) in single bonds	about 150–250 nm
π → π*	double and triple bonds, aromatic rings	about 170–300 nm for isolated bonds; longer with conjugation
n → π*	lone pair next to a double bond, e.g. C=O	about 270–350 nm, weak
d → d	transition metal ions	visible region
charge transfer	complexes such as MnO₄⁻	visible, very intense

Groups that absorb in the UV–visible region are called chromophores: C=C, C=O, N=N, aromatic rings and conjugated systems.

Conjugation: the key to colour

A conjugated system is a chain of alternating single and double bonds, in which electrons are delocalised over many atoms. The more extended the conjugation:

• the smaller the energy gap between the highest occupied and lowest unoccupied orbitals
• the longer the wavelength of light absorbed
• and (often) the stronger the absorption

Molecule	Conjugated C=C bonds	Approximate λmax
Ethene	1	170 nm (UV)
Buta-1,3-diene	2	217 nm (UV)
Hexa-1,3,5-triene	3	258 nm (UV)
β-carotene	11	about 450 nm (visible: absorbs blue)

β-carotene absorbs blue light, so it looks orange. That’s why carrots are orange. Potato starch has no extended conjugation, so it absorbs no visible light and looks white.

The same principle explains many dyes, indicators and pigments. Phenolphthalein is colourless in acid but pink in alkali because losing protons links its rings into one large conjugated system. See phenolphthalein.

Colour and complementary colours

A substance looks coloured because it absorbs some wavelengths of white light and transmits or reflects the rest. The colour we see is roughly the complementary colour of the light absorbed:

Wavelength absorbed (nm)	Colour absorbed	Colour seen
400–430	violet	yellow-green
430–490	blue	orange-yellow
490–560	green	red-purple
560–590	yellow	violet-blue
590–650	orange	blue
650–700	red	green-blue

Copper(II) sulfate solution absorbs mainly orange-red light (around 800 nm and extending into the red), so it looks blue.

Transition metal colours

Transition metal ions often have partly filled d orbitals. In a complex, the ligands split the d orbitals into two groups with slightly different energies. An electron can jump between them by absorbing visible light: a d–d transition.

The size of the split, and so the colour, depends on:

• the metal
• its oxidation state
• the ligands
• the shape of the complex

That’s why [Cu(H₂O)₆]²⁺ is pale blue but [Cu(NH₃)₄(H₂O)₂]²⁺ is deep blue-violet: ammonia splits the d orbitals more strongly, shifting the absorption to shorter wavelength and increasing its intensity.

d–d transitions are relatively weak (small molar absorptivity), which is why many transition metal solutions are pale. Charge-transfer transitions, like those in the permanganate ion, are much stronger, giving intense colours at low concentrations.

The absorption spectrum

A UV–visible spectrum is a graph of absorbance (y-axis) against wavelength (x-axis). Peaks are broad, because each electronic transition is accompanied by many vibrational and rotational changes.

The two most useful features:

• λmax: the wavelength of maximum absorbance. It helps identify a chromophore and is the best wavelength for quantitative measurements.
• Absorbance at λmax: related to concentration by the Beer–Lambert law, A = εcl.

The spectrophotometer

A UV–visible spectrophotometer has:

1. Light sources: a deuterium lamp for UV and a tungsten–halogen lamp for visible light.
2. Monochromator: a diffraction grating that selects a narrow band of wavelengths.
3. Sample holder: a cuvette with a 1 cm path length. Quartz cuvettes are used for UV, because ordinary glass and most plastics absorb UV below about 300–350 nm.
4. Detector: a photodiode or photomultiplier measuring the transmitted light.

Double-beam instruments split the light into two beams, one through the sample and one through a reference (blank), and compare them continuously. Diode-array instruments record the whole spectrum at once.

Uses of UV–visible spectroscopy

• Measuring concentration of coloured or UV-absorbing substances using calibration curves.
• Biochemistry: DNA absorbs strongly at 260 nm and proteins at 280 nm (due to aromatic amino acids). The ratio A₂₆₀/A₂₈₀ checks the purity of DNA samples.
• Kinetics: following the rate of a reaction by monitoring how a reactant’s or product’s absorbance changes.
• Pharmaceuticals: assays of drug content and dissolution testing.
• Water analysis: nitrate, phosphate and metals (after adding colour-forming reagents).
• HPLC detectors: most HPLC systems use UV absorbance to detect compounds as they elute. See HPLC.
• Sunscreens: UV filters are chosen for strong absorption in the UVA (315–400 nm) and UVB (280–315 nm) ranges.

Worked example: why do carrots and tomatoes differ?

Both carrots and tomatoes get their colour from long conjugated molecules. β-carotene in carrots has 11 conjugated C=C bonds and absorbs most strongly around 450 nm (blue), so it looks orange. Lycopene in tomatoes has 11 conjugated C=C bonds plus two extra isolated ones, and a slightly different shape that extends its absorption further towards green light, so tomatoes look red. Small changes in conjugation make clearly visible differences in colour.

Limitations

• Spectra are broad and give limited structural information compared with IR or NMR.
• Only molecules with suitable chromophores absorb in the accessible range; many saturated compounds are “invisible”.
• The Beer–Lambert law only holds for fairly dilute solutions.
• Solvents must be transparent in the region measured; ethanol, water and hexane are common choices for UV work.

Key takeaways

• UV–visible spectroscopy measures absorption at about 200–800 nm, caused by electrons moving to higher energy levels.
• Chromophores such as C=C, C=O and aromatic rings absorb; more conjugation shifts absorption to longer wavelengths.
• We see the complementary colour of the light absorbed.
• Transition metal colours come from d–d transitions, which depend on the metal, oxidation state and ligands.
• The main uses are measuring concentration (Beer–Lambert law) and monitoring reactions.