# What causes different colors of light photons to be emitted?

Light is a form of electromagnetic radiation that is visible to the human eye. The color we perceive light to be depends on the wavelength (or frequency) of the photons that make up that light. Different wavelengths of light appear as different colors to us. This article will examine what causes photons of light to have different wavelengths and therefore appear as different colors.

Light is composed of discrete particles called photons. Photons are the smallest units or “packets” of light. They have properties of both waves and particles. As particles, they travel in a straight line and have energy and momentum. As waves, they oscillate with a certain frequency and wavelength.

The wavelength of a photon corresponds to the color of light we see. The visible spectrum of light that humans can see ranges from about 380 nanometers (violet) to about 740 nanometers (red). Other wavelengths exist outside the visible spectrum, such as radio waves, microwaves, infrared radiation, ultraviolet radiation, X-rays and gamma rays. These are all forms of electromagnetic radiation that differ only by their wavelength.

## What Determines a Photon’s Wavelength?

A key factor that determines a photon’s wavelength is the amount of energy it contains. Photons with higher energy have shorter wavelengths, while lower energy photons have longer wavelengths.

The energy of a photon is directly proportional to its frequency and inversely proportional to its wavelength. This relationship is described by Planck’s equation:

E = hf = \frac{hc}{\lambda}

Where:

• E = Energy of the photon (Joules)
• h = Planck’s constant (6.626 x 10-34 Joules-second)
• f = Frequency of the photon (Hertz)
• c = Speed of light in a vacuum (2.998 x 108 meters/second)
• \lambda = Wavelength of the photon (meters)

So photons with higher frequencies have more energy and shorter wavelengths. Lower frequency photons have less energy and longer wavelengths.

## Energy Levels in Atoms

What determines the specific energy and frequency of a photon? The answer lies within atoms.

Atoms can only exist in certain allowed energy levels or states. Electrons bound to atoms can transition between these discrete energy levels by absorbing or emitting photons.

An electron can jump from a lower energy level to a higher one by absorbing a photon with energy equal to the exact difference between those levels.

Conversely, electrons can also drop to lower energy levels, emitting a photon in the process. The energy of the emitted photon corresponds exactly to the difference in energy between the levels.

## Emission Spectra

The specific collection of photon energies an atom can emit is unique to each element, creating a distinctive emission spectrum.

When heated, different atoms emit different colors of light based on these emission spectra. For example:

• Hydrogen atoms emit red light
• Sodium atoms emit yellow light
• Mercury atoms emit blue light

This is why neon signs containing different gases glow with different colors. The color emitted depends on the energies required for electrons in that element to transition between energy levels.

## The Electromagnetic Spectrum

Not all electron energy transitions in atoms produce photons in the visible range. Some photons have energies corresponding to ultraviolet or infrared wavelengths.

Ordered by increasing photon energy and decreasing wavelength, the electromagnetic spectrum categorizes all types of electromagnetic radiation.

Type of Radiation Wavelength Range Photon Energy Range
Radio waves 1000 mm – 1 mm 1.24 x 10-6 eV – 1.24 x 10-3 eV
Microwaves 100 mm – 1 mm 1.24 x 10-3 eV – 1.24 eV
Infrared 750 nm – 1 mm 1.24 eV – 1.7 eV
Visible light 390 nm – 750 nm 1.7 eV – 3.1 eV
Ultraviolet 10 nm – 390 nm 3.1 eV – 124 eV
X-rays 0.01 nm – 10 nm 124 eV – 124 keV
Gamma rays > 124 keV

As shown here, visible light makes up only a small slice of the full electromagnetic spectrum. But it is the range of photon energies our eyes can detect as color.

## Discrete Nature of Photons

It’s important to understand that photons themselves have discrete amounts of energy, set by their frequency according to Planck’s equation. They can only be emitted or absorbed by electrons in discrete jumps between allowed energy levels in atoms and molecules.

This quantization of photon energy into distinct packets is a key aspect of the quantum mechanical nature of light. While light can act as a wave, it is fundamentally composed of these indivisible photon particles.

The set energy levels in atoms and quantized energies of photons they emit gives rise to distinct spectral emission lines. Exciting an atom with a range of energies still only produces photons at specific wavelengths corresponding to allowed transitions.

## Photon Interactions

When a photon is absorbed by matter, its energy can be transferred in different ways:

• Photochemical interactions – The photon energy induces a chemical change by breaking or altering molecular bonds.
• Photoelectric effect – The photon ejects electrons from a surface by transferring all its energy to kinetic energy of electrons.
• Compton scattering – The photon scatters off an electron, losing energy while transferring momentum to the electron.
• Pair production – Extremely energetic photons convert into matter in the form of an electron-positron pair.

The way photons interact depends on both their energy, the medium, and how susceptible electrons are to excitation. Visible light photons mainly cause electrons to transition between energy levels or molecular vibrations and rotations. Higher energy photons can eject, scatter, or produce electrons.

## Conclusion

In summary, light comes in discrete packets called photons that have wave-particle duality. A photon’s wavelength determines what color we perceive it to be. The photon’s wavelength and frequency are intrinsically linked to the energy it carries.

Electrons in atoms can only exist at certain allowed energy levels unique to each element. Transitions between these levels result in emission or absorption of photons at energies corresponding to the exact change in energy. The set of photon energies an atom emits gives each element a distinctive set of spectral lines.

So in the end, the fundamental reason light comes in different colors is due to the quantized nature of photons and electron energy levels in atoms that determine the allowed energies and frequencies of light that can be emitted or absorbed. Understanding these quantum mechanical processes helped advance modern physics and our comprehension of light beyond classical explanations.