Sunlight excites the electrons in a molecule. Specifically, when a photon of the right energy strikes a molecule, it bumps one of the molecule’s electrons from a lower-energy orbital to a higher-energy one. This jump is called an electronic transition, and it’s the fundamental event behind everything from photosynthesis to sunburn to human vision.
Electrons Jump Between Orbitals
Electrons in a molecule don’t float around randomly. They occupy specific energy levels called molecular orbitals. In the ground state (the molecule’s resting condition), electrons fill the lowest available orbitals first. The highest orbital that contains an electron is called the HOMO, and the lowest empty orbital waiting above it is called the LUMO. When sunlight delivers a photon whose energy exactly matches the gap between these two orbitals, an electron absorbs that energy and jumps from the HOMO up to the LUMO.
Different types of electrons make different types of jumps. In a simple single bond, an electron in a bonding orbital can be promoted to an antibonding orbital. In molecules with double bonds, electrons in the pi bond jump to a higher-energy pi antibonding orbital. Molecules with atoms like oxygen or nitrogen that carry lone pairs of electrons can undergo yet another type of transition, where one of those lone-pair electrons gets promoted to an antibonding orbital. These distinctions matter because each type of jump requires a different amount of energy, which determines what wavelength of light the molecule absorbs.
Different Wavelengths Cause Different Effects
Sunlight contains ultraviolet, visible, and infrared radiation, and each portion interacts with molecules differently. UV and visible light (roughly 200 to 700 nanometers) carry enough energy, between 36 and 143 kilocalories per mole, to push electrons into higher orbitals. This is why absorption in these wavelength ranges is sometimes called “electronic spectroscopy.” Infrared light, by contrast, is lower in energy. It doesn’t excite electrons at all. Instead, it makes the bonds within a molecule vibrate faster, which is why infrared radiation feels warm on your skin.
Within the UV and visible range, the exact wavelength a molecule absorbs depends on the size of its HOMO-LUMO gap. A large gap means the molecule needs high-energy UV photons. A small gap means lower-energy visible light will do the job, which is why many colored substances (dyes, pigments, autumn leaves) have electrons that are easy to excite with visible wavelengths.
Chromophores: Where Absorption Happens
Not every part of a large molecule responds to sunlight equally. The specific region where light absorption occurs is called a chromophore. A chromophore is typically an unsaturated functional group, meaning it contains double bonds or systems of alternating single and double bonds (conjugated systems). The more extended the conjugation, the smaller the energy gap between orbitals, and the longer the wavelength of light the molecule can absorb. This is why many biological pigments and synthetic dyes are built around long chains or rings of conjugated double bonds.
Neighboring groups on the molecule can shift absorption slightly by donating or withdrawing electron density from the chromophore, but the chromophore itself is where the electronic transition takes place.
How This Powers Photosynthesis
Chlorophyll is a perfect case study. The part of chlorophyll that captures sunlight is its porphyrin ring, a large flat ring of conjugated double bonds with a magnesium atom at the center. When a photon hits this ring, it promotes an electron from the highest occupied orbital into the lowest unoccupied orbital. The dominant transition in chlorophyll involves exactly this HOMO-to-LUMO jump.
Interestingly, the protein environment around chlorophyll slightly bends and distorts the porphyrin ring, which shrinks the HOMO-LUMO gap and shifts absorption toward longer (redder) wavelengths by up to 17 nanometers. This tuning allows different chlorophyll molecules within the same photosynthetic complex to absorb slightly different colors of light, broadening the range of sunlight a plant can harvest.
How This Enables Vision
Your ability to see depends on the same electron-excitation principle, but with a twist. Inside the photoreceptor cells of your retina, a small molecule called retinal sits covalently attached to a protein called opsin. Retinal’s chromophore is a chain of conjugated double bonds. When a photon excites the electrons in this chain, the added energy doesn’t just bump an electron to a higher orbital; it causes the molecule to physically change shape. A specific double bond in retinal (at the 11th carbon position) flips from a bent “cis” configuration to a straight “trans” configuration. This shape change triggers the opsin protein to activate, launching the cascade of signals your brain interprets as sight.
This photoisomerization, where light energy rearranges the geometry around a double bond, is one of the fastest and most efficient reactions in biology.
How This Causes DNA Damage
The same mechanism that powers vision and photosynthesis can also cause harm. DNA bases, particularly thymine and cytosine, contain double bonds that absorb UV light. When UV photons excite the electrons in these double bonds, the added energy can open the bond and allow it to react with a neighboring base on the same DNA strand. If two thymine bases sit next to each other, this can create a thymine dimer: two covalent bonds locking the adjacent bases together. These dimers distort the DNA helix and, if not repaired, can lead to mutations and skin cancer.
DNA has some built-in resilience, though. When UV light excites the electrons in pyrimidine bases (thymine, cytosine, uracil), the molecule usually converts that electronic energy into heat within picoseconds, far faster than a chemical reaction can occur. The excited electron drops back to the ground state through rapid nonradiative decay, and the photon’s energy dissipates as vibrations in the surrounding water molecules. Only a small fraction of excitation events lead to actual dimer formation.
How Sunscreen Exploits This Process
Chemical sunscreens work by containing molecules whose electrons are deliberately easy to excite with UV light. Oxybenzone, one common UV-absorbing ingredient, absorbs a UV photon and undergoes an ultrafast internal rearrangement: the excited molecule briefly shifts into a different structural form, then snaps back to its original shape while dumping the absorbed energy as harmless heat. This entire cycle, absorption, rearrangement, and energy release, happens so quickly that there’s no time for the energy to cause damage to skin cells underneath.
In every one of these examples, from plant pigments to your retina to sunscreen, the story starts in the same place: a photon of sunlight delivers its energy to an electron, and that electron jumps to a higher orbital. What happens next depends on the molecule, but the initial event is always an electronic transition.