A photon is a single packet of light energy, the smallest possible unit of electromagnetic radiation. In chemistry, photons matter because they are the mechanism by which light interacts with atoms and molecules, driving everything from the color of a solution to photosynthesis to the way chemists identify unknown compounds. Every time light breaks a bond, excites an electron, or gets absorbed by a molecule, it does so one photon at a time.
The Energy of a Single Photon
The defining equation in all of photon chemistry is deceptively simple: E = hν. The energy (E) of a photon equals Planck’s constant (h) multiplied by the frequency (ν) of the light. Planck’s constant is an extremely small number, 6.626 × 10⁻³⁴ joule-seconds, which tells you that a single photon carries a tiny amount of energy. But that tiny amount is fixed for any given frequency of light. A photon of red light always carries the same energy as any other photon of red light.
Because frequency and wavelength are inversely related, you can also write this as E = hc/λ, where c is the speed of light and λ is the wavelength. Shorter wavelengths mean higher energy. This is why ultraviolet light can damage DNA while infrared light just warms your skin. The photons in UV light carry enough energy per packet to break chemical bonds, while infrared photons only have enough energy to make molecules vibrate faster.
The idea that energy comes in these discrete packets, rather than as a continuous stream, was revolutionary. Max Planck proposed the concept in 1900 to solve a problem in physics, and Einstein extended it in 1905 by arguing that light itself is made of these packets, which he called photons. This insight earned Einstein the Nobel Prize and laid the foundation for quantum mechanics.
How Photons Interact With Atoms and Molecules
Electrons in atoms and molecules sit at specific energy levels. They can’t exist at energies between those levels. When a photon hits a molecule and its energy exactly matches the gap between two electron energy levels, the molecule absorbs the photon and the electron jumps to the higher level. If the photon’s energy doesn’t match any available gap, it passes through or bounces off. This selectivity is why different substances absorb different colors of light.
Once an electron has been excited to a higher energy level, it doesn’t stay there long. It eventually falls back down, and when it does, it releases a photon. This is the basis of two phenomena you may have encountered: fluorescence and phosphorescence. In fluorescence, the electron drops back quickly, emitting light almost instantly. In phosphorescence, the electron gets temporarily trapped in an intermediate state and takes much longer to release its photon, sometimes seconds. This is why glow-in-the-dark materials keep emitting light after you turn off the lamp.
Different Photon Energies, Different Chemical Effects
The electromagnetic spectrum is really a spectrum of photon energies, and each energy range does something different to molecules:
- Infrared photons carry just enough energy to make chemical bonds vibrate and rotate. This is why infrared light feels warm on your skin: it increases the kinetic energy of your molecules. Chemists exploit this in infrared spectroscopy, where the specific vibrations a molecule absorbs reveal which types of bonds it contains.
- Visible and ultraviolet photons carry enough energy to promote electrons from bonding orbitals to higher-energy orbitals. This is what gives colored compounds their color: they absorb certain visible wavelengths and reflect the rest. UV photons are energetic enough to actually break certain bonds. For example, the energy needed to break a hydrogen-hydrogen bond corresponds to a photon with a wavelength around 280 nm, which falls in the far ultraviolet range.
- X-ray photons are so energetic they can knock electrons completely out of atoms, ionizing them. This is why X-ray exposure requires safety precautions.
A practical example of UV damage: when a UV photon strikes a DNA molecule, it can be absorbed by an electron in one of the bonds within a thymine base. The photon promotes that electron into an antibonding orbital, effectively breaking the bond and creating a highly reactive site. This is the molecular mechanism behind sunburn and UV-induced mutations.
Photons Driving Chemical Reactions
In thermal chemistry, reactions need activation energy, and that energy comes from heat. Molecules crash into each other with enough force to break and rearrange bonds. Photochemistry works differently. A photon lifts a molecule directly into a high-energy excited state, bypassing the need to climb the energy barrier from below. As one description puts it, the molecule comes down from above rather than confronting a difficult ascending path from the ground state.
This opens up reactions that would be impossible or impractical with heat alone. Carbon-carbon bonds, for instance, can form under light irradiation without requiring the chemical bases that thermal reactions need. The absorbed light energy breaks existing bonds, generating high-energy intermediates that then release their energy by forming new bonds and yielding products.
Chemists measure the efficiency of these light-driven reactions using a value called quantum yield. Defined by IUPAC as the number of defined events occurring per photon absorbed, quantum yield tells you how many molecules react for each photon the system absorbs. A quantum yield of 1.0 means every absorbed photon produces one reaction event. Values above 1.0 are possible when one photon triggers a chain reaction.
Photosynthesis: Photons Building Molecules
The most important photochemical reaction on Earth is photosynthesis. Chlorophyll molecules absorb photons in the red and blue parts of the visible spectrum and reflect green, which is why plants look green. Inside plant cells, specialized chlorophyll pairs absorb red light at very specific wavelengths. One pair, called P680, absorbs most efficiently at 680 nm. Another, P700, absorbs at 700 nm.
When one of these chlorophylls absorbs a photon, the energy excites an electron to a higher energy level. That energized electron then passes through a chain of acceptor molecules, and the energy it releases along the way is used to split water into oxygen, protons, and electrons. The end products are the energy-storage molecules ATP and NADPH, which the plant then uses to build sugars from carbon dioxide. Every oxygen molecule you breathe exists because a photon kicked an electron loose from a chlorophyll molecule.
How Chemists Use Photons in the Lab
Spectroscopy, one of the most widely used analytical techniques in chemistry, is built entirely on how molecules absorb photons. In UV-visible spectroscopy, a beam of light passes through a sample solution. The instrument scans through wavelengths and records which ones the sample absorbs and how strongly. The resulting absorption spectrum acts like a molecular fingerprint.
The relationship between photon absorption and concentration follows the Beer-Lambert Law: A = εlc. Absorbance (A) equals the molar absorptivity (ε) times the path length (l) times the concentration (c). Molar absorptivity is a constant specific to each substance at each wavelength, measured in L mol⁻¹ cm⁻¹. It reflects how strongly that substance absorbs photons at that particular energy. For example, guanosine (a building block of DNA) has a molar absorptivity of 8,400 at its peak absorption wavelength of 275 nm.
This means that if you know what a substance is, you can measure its concentration by shining light through a sample and reading the absorbance. Or, if you have an unknown substance, its absorption pattern can help you identify it. The whole method works because photon absorption is quantized: specific molecules absorb specific wavelengths in predictable, reproducible ways.
Wave-Particle Duality
One of the stranger aspects of photons is that they behave as both waves and particles, depending on how you observe them. The photoelectric effect, where light striking a metal surface ejects electrons, only makes sense if light is made of particles. Each photon delivers a discrete punch of energy to a single electron. But interference and diffraction patterns, where light bends around edges and creates alternating bright and dark bands, only make sense if light is a wave.
In 1924, Louis de Broglie proposed that this duality isn’t unique to photons. All moving objects have wave-like properties. For anything with significant mass, the wavelength is far too small to detect. But for photons, which have no rest mass, the wave nature is just as apparent as the particle nature. In chemistry, both aspects matter. The wave nature determines a photon’s wavelength and frequency, which define what kind of molecular interaction it can have. The particle nature means that energy transfers happen in discrete, all-or-nothing events: an electron either absorbs a whole photon or none of it.