The visible region of the electromagnetic spectrum lies between the ultraviolet (UV) and the infrared (IR) regions, and spans wavelengths of approximately 400-800 nm. The shorter wavelengths (violet, blue, and green) are higher in energy and border the UV region, whereas the longer wavelengths (yellow, orange and red) are relatively lower in energy and border the IR region. The colors progress from short to long wavelengths in a reverse rainbow, Violet -> Blue -> Green -> Yellow -> Orange -> Red, as shown below, and a color wheel can be created by putting these colors consecutively around a pie (see below).

When a compound absorbs light in this region, it can promote an electron from the highest occupied molecular orbital (HOMO) to the lowest unoccupied molecular orbital (LUMO) which has higher energy. This energy difference between orbitals, ΔE, corresponds to the energy or wavelength of light required to promote an electron from the HOMO to the LUMO and these two are related by the equation, E = hc/λ, where E = energy (in Joules, J), h = Planck’s constant (6.626×10-34 J·s), c = speed of light (~3×10⁸ m/s), and λ = wavelength (m). Therefore, when the energy difference between orbitals is greater (large HOMO-LUMO gap), the wavelength of light required to promote an electron, λ, is small (short wavelength, high energy light, i.e. violet or blue), and vice versa. In general, extended conjugated systems have smaller HOMO-LUMO gaps and therefore the wavelength of light they can absorb to promote an electron is larger/lower energy (i.e. red or orange). In fact, the molecular structures of dye molecules are specifically designed to have a certain extent of pi conjugation in order to absorb in a certain area of the visible region.

A visible absorption spectrum of a compound, like shown below, plots the absorbance of the sample on the y-axis, vs. the wavelengths of light (in nm) on the x-axis. If a compound absorbs a certain wavelength of light which has enough energy to promote an electron to a higher energy orbital, there will be a peak in the spectrum, labeled with the highest point as, λ_max. The wavelengths of absorbance are related to the color of the compound, as anything we observe to be colored appears to be a certain color based on which wavelengths of light are absorbed vs. reflected. So, a compound that appears colored to us, does so because it absorbs the wavelengths in the visible region of the opposite, or complementary color on a color wheel. For example, the compound which gives the spectrum below when visible light is passed through absorbs light in the blue-green region, therefore, looking across the color wheel, it will appear red-orange.

Indicators are substances whose solutions change color due to changes in pH. These are called acid-base indicators. They are usually weak acids or bases, but their conjugate base or acid forms have different colors due to differences in their absorption spectra. Indicators are organic weak acids or bases with complicated structures. We represent a general indicator by the formula HIn. H-indicator absorbs at a certain wavelength and is of one color. Indicator- absorbs at a different length, so it is a different color. At a low pH, there is a high concentration of hydrogen ions, so H-indicator and its color will predominate. At a high pH, there is a low concentration of hydrogen ions, so the indicator and its color will predominate. At a neutral pH both H-indicator and indicator^– will co-exist in equilibrium, so there will be a mixture of the two.

The colors of the universal indicator are as follows:

• Red: Very Acidic
• Orange: Acidic
• Yellow: Weakly acidic
• Green: Neutral
• Blue: Basic
• Purple: Very Basic