A chromophore is the molecule within an object that gives it color. It does this by absorbing some wavelengths of light present in the visible spectrum, while reflecting others. The discs inside the outer segments of photoreceptors are lined with a vast amount of chromophore molecules, bound to opsin proteins. These chromophores are called retinal. They give photoreceptors the ability to respond to photons of light.

To understand the principles of how a chromophore absorbs light, an idea of molecular orbital theory, as well as its underlying concept of molecular orbitals, is needed. This is a chemical theory which proposes a mechanism for the manner in which atoms bind to each other. In it, atoms form associations in the process of becoming molecules using either σ (sigma) or 𝛑 (pi) bonds.

Classic sigma and pi bond image

When two atomic orbitals combine, the same number of molecular orbitals are created. These two molecular orbitals are separated into a bonding and an antibonding orbital. If an energy diagram of a molecule is drawn, the least stable (highest in energy) bonding orbital with an electron inside it is called the HOMO - the highest occupied molecular orbital. The antibonding orbital just above it in terms of energy, empty of electrons, is called the LUMO - the lowest unoccupied molecular orbital.

There exists an energy difference between the HOMO and the LUMO, which is important because energy and photons are concepts directly connected by the following relationship:

λ = h*c/ΔE

This relationship makes clear that the shorter a wavelength is, the more energy it carries.

A photon is absorbed when it gives up its energy to a molecule which expends it by promoting an electron from the HOMO to the LUMO. The larger the energy gap is, the shorter the wavelength of light characterizing the photon must be. This mechanism underlies light absorption by chromophores.

How exactly does this process occur in chromophores, like the photoreceptor molecule retinal? When two adjacent atomic orbitals known as p-orbitals interact, they form 𝛑 bonds, which are more widely known as ‘double bonds’. However, if a molecule has three or more p-orbitals positioned adjacent to each other, a conjugated 𝛑-system may result. The formation of this 𝛑-system requires that p-orbitals overlap, so the series of atoms comprising it must have p-orbitals oriented in a planar fashion.

A conjugated 𝛑-system lowers the overall energy of a molecule, which increases its stability. Conventionally such a system is drawn as below, with alternating single and double bonds along the backbone of the molecule.

A conjugated system such as this allows for the delocalization of the electrons it contains across all of the p-orbitals comprising it. Put differently, the region of probability density of these electrons is not restricted to just one atom, but rather to a long series of atoms. The electrons of each orbital are ‘shared’ by the entire system.

The 𝛑-systems described above occur in molecules that have atoms engaging in sp or sp^2 hybridization. These types of hybridization provide the necessary unhybridized p-orbitals, the lobes of which can overlap if they’re aligned on the same plane. Because an sp^3 hybridized atom does not have a free p-orbital, its presence terminates the sequence of overlapping p-orbitals, and thus also terminates the 𝛑-system.

A p-orbital which forms a link in a 𝛑-system may contribute one electron, two electrons (in which case it is known as a ‘lone pair’) or no electrons at all (an ‘empty’ orbital).

Chromophores employ such conjugated 𝛑-systems to capture photons of light. Within these systems, 𝛑 electrons move from the ‘highest occupied molecular orbital’ (HOMO) to the ‘lowest unoccupied molecular orbital’ (LUMO) if a photon of the right wavelength imparts the right right amount of energy to ‘bridge’ this gap.

To restate what I’ve mentioned above, the relationship between energy and wavelength is λ = h*c/ΔE. This means that shorter wavelengths carry more energy. The energy difference between the HOMO and the LUMO sets the wavelength that a 𝛑-system will absorb. The longer the system is, the longer the wavelength it will absorb.

Very short systems of conjugated p-orbitals have a large energy gap between the HOMO and the LUMO - thus, they can only absorb very short, energetic wavelengths. Because these wavelengths occur in the UV spectrum, to the human eye such an object will appear colorless.

Longer systems will appear progressively ‘bluer’ - shifting from red, to yellow, to green, to blue. Pigments that appear green or blue cannot rely solely on such conjugated 𝛑-systems to lower the energy between the HOMO and the LUMO. The rarity of the color blue in nature is a testament to how difficult it is to sufficiently lower this transition energy.