Ultraviolet Radiation

Ultraviolet (UV) radiation is electromagnetic radiation with wavelengths ranging from about 10 nm to 400 nm — just shorter than visible violet light (around 400 nm) and longer than X-rays. Because wavelength and frequency are inversely related by c = fλ, UV photons vibrate at frequencies of roughly 7.5 × 10¹⁴ Hz to 3 × 10¹⁶ Hz. That places them firmly in the high-energy portion of the electromagnetic spectrum, energetic enough to break certain chemical bonds and ionise atoms. The Sun is Earth's primary UV source, but the atmosphere filters most of it. The ozone layer (O₃ molecules in the stratosphere, 15–35 km altitude) absorbs almost all UV-C (10–280 nm) and most UV-B (280–315 nm), allowing mainly UV-A (315–400 nm) to reach the surface. This filtering is critical to life: unfiltered UV-C and UV-B would denature proteins and shred DNA in surface organisms within minutes. Artificial UV sources include mercury-vapour lamps, tanning beds, germicidal sterilisers, and blacklights. UV radiation has enormous practical importance. On the beneficial side, UV-B triggers vitamin D synthesis in human skin, UV light sterilises hospital equipment by destroying microbial DNA, and astronomers use UV telescopes to study hot stars and galaxies. On the harmful side, UV-B and UV-A damage skin DNA, leading to sunburn, premature ageing, and increased skin cancer risk. Understanding UV requires mastering how photon energy relates to frequency and wavelength — which is exactly where the physics begins.

Key takeaways
  • Position on the Electromagnetic Spectrum
  • Photon Energy and the Planck Relation
  • UV Sub-Bands and Their Effects
  • DNA Damage: The Molecular Mechanism
  • Ozone Absorption and the UV Index

Position on the Electromagnetic Spectrum

All electromagnetic waves travel at c ≈ 3 × 10⁸ m/s in a vacuum. They are ordered by wavelength: radio → microwave → infrared → visible → ultraviolet → X-ray → gamma. UV sits immediately above visible light. Its wavelengths (10–400 nm) are shorter than any colour your eye detects, so UV is literally 'beyond violet'. Because f = c/λ, shorter wavelength means higher frequency — and higher energy per photon.

Photon Energy and the Planck Relation

The energy carried by a single UV photon is given by Planck's relation: E = hf = hc/λ. Using h = 6.626 × 10⁻³⁴ J·s, a UV-A photon at λ = 380 nm has E ≈ 3.3 eV, while a UV-C photon at λ = 200 nm has E ≈ 6.2 eV. For comparison, visible red light at 700 nm gives only 1.8 eV. Most organic chemical bonds (C–H, C–C, O–H) have bond energies in the range 3–5 eV, which explains why UV-B and UV-C can break them while visible light cannot. This is the root of all UV photochemistry.

UV Sub-Bands and Their Effects

Physicists and health authorities divide UV into three sub-bands. UV-A (315–400 nm, ~3.1–3.9 eV): penetrates deep into skin, causes indirect DNA damage through reactive oxygen species, responsible for tanning and long-term ageing. UV-B (280–315 nm, ~3.9–4.4 eV): directly absorbed by DNA bases; causes pyrimidine dimer formation (adjacent thymine bases bond abnormally), leading to sunburn and mutation. UV-C (10–280 nm, ~4.4–124 eV): most destructive, almost entirely blocked by ozone and oxygen; used in germicidal lamps because it is lethal to bacteria and viruses.

DNA Damage: The Molecular Mechanism

DNA absorbs UV-B strongly at ~260–270 nm because the aromatic rings in thymine and cytosine bases resonate at those frequencies. When two adjacent thymine bases on the same DNA strand absorb UV photons, a [2+2] cycloaddition reaction fuses their carbon–carbon double bonds, creating a thymine dimer. This kink in the DNA helix stalls replication and transcription. Cells have repair mechanisms (nucleotide excision repair), but repeated or intense UV exposure overwhelms repair, leading to mutations in tumour-suppressor genes like p53 — a key step toward melanoma. This is quantitative chemistry driven by photon energy exceeding bond thresholds.

Ozone Absorption and the UV Index

Ozone absorbs UV via the Hartley band (200–310 nm) and Huggins band (310–340 nm). The Beer–Lambert law describes how intensity decreases through the ozone layer: I = I₀ e^(−σ n L), where σ is the absorption cross-section of O₃, n is the number density of ozone molecules, and L is the path length through the layer. The UV Index (UVI) is a standardised scale (0–11+) reporting the erythemally weighted irradiance at the surface — it accounts for the fact that UV-B is far more biologically damaging per photon than UV-A.

The Key Equation

E=hcλE = \frac{hc}{\lambda}
  • E — Photon energy (Joules (J) or electron-volts (eV)): The energy carried by one quantum of electromagnetic radiation. 1 eV = 1.602 × 10⁻¹⁹ J.
  • h — Planck's constant (J·s): h = 6.626 × 10⁻³⁴ J·s; the fundamental quantum of action relating energy to frequency.
  • c — Speed of light in vacuum (m/s): c = 2.998 × 10⁸ m/s; the universal speed limit for electromagnetic waves.
  • \lambda — Wavelength (metres (m); often nanometres (nm) for UV): The distance between successive wave crests. UV spans 10–400 nm. Remember: 1 nm = 10⁻⁹ m.

Worked Example

Problem. A UV-B photon has a wavelength of 300 nm. (a) Calculate its energy in joules. (b) Convert to electron-volts. (c) The bond dissociation energy of a C–C bond is about 3.6 eV. Can this photon break a C–C bond?

E=hcλ≈6.62×10−19 J≈4.13 eVE = \frac{hc}{\lambda} \approx 6.62 \times 10^{-19}\,\text{J} \approx 4.13\,\text{eV}

Answer. E ≈ 6.62 × 10⁻¹⁹ J ≈ 4.13 eV; yes, it can break a C–C bond.

Figure
Diagram
Ultraviolet Radiation: Spectrum and Sub-bandsVisible LightUV-AUV-BUV-C700 nm400 nm315 nm280 nm10 nmIncreasing frequency & photon energy →Photon Energy (E = hc/λ)UV-A: 3.1–3.9 eVTanning, ageingUV-B: 3.9–4.4 eVDNA dimers, sunburnUV-C: 4.4–124 eVGermicidal; blocked by O₃UV-B Thymine Dimer FormationTTNormal DNA→UV-BphotonT—TdimerDamaged DNA
Origins · 1 minJohann Wilhelm Ritter and how this was discovered
JW
Johann Wilhelm Ritter
1776–1810 · German

Ritter discovered ultraviolet radiation in 1801, just one year after William Herschel discovered infrared. He noticed that silver chloride (AgCl) — which darkens when exposed to light — was darkened most powerfully by invisible radiation just beyond the violet end of the spectrum. He called it 'chemical rays' (chemische Strahlen), recognising that UV had a strong chemical effect distinct from heat. His systematic experiments established that the invisible rays beyond violet behaved like light (they reflected and refracted) but were far more chemically active — the first experimental evidence linking high-frequency EM radiation to chemical reactions. Ritter died at only 33 years old, likely from exhaustion due to overwork. A brilliant but troubled polymath, he also experimented with electrochemistry and once performed electrical self-experiments that severely damaged his health.

  • X-Ray Radiation — X-rays are the next band beyond UV-C; understanding how photon energy scales with frequency across these bands shows why X-rays are even more penetrating and ionising than UV.
  • The Visible Light Spectrum — Visible light is immediately adjacent to UV-A; comparing photon energies (1.8–3.1 eV for visible vs 3.1+ eV for UV) explains why UV crosses the threshold for photochemical damage.
  • Planck's Relation: E = hf — The Planck relation is the central equation for calculating UV photon energies and understanding why higher frequency means more energetic — and more damaging — radiation.
  • Wave Speed, Frequency, and Wavelength — The relationship c = fλ is essential for converting between UV wavelengths (given in nm) and frequencies, and for locating UV on the electromagnetic spectrum.