Absorption and Emission of Light

The Bohr model describes electrons as occupying specific, quantized energy levels — or orbits — around the nucleus. Because these levels are quantized, electrons cannot exist between them. An electron can only move from one energy level to another by absorbing or emitting energy.

Key Takeaways

• The Bohr model treats electron energy levels as quantized — transitions between them happen only via absorption or emission of a photon matching the energy gap exactly.

• Absorption moves an electron to a higher level (excited state); emission releases a photon as an electron falls to a lower level. Both follow E = hf.

• Absorption spectroscopy (including IR and UV-Vis) measures wavelengths absorbed; emission spectroscopy measures wavelengths released, as seen in flame-test spectral lines.

• Fluorescence: high-energy light absorbed, lower-energy visible light re-emitted, with the wavelength shift known as the Stokes shift.

Absorption: Moving to a Higher Energy Level

When an electron absorbs energy, it jumps to a higher energy level, entering an excited state. That energy comes from a photon whose energy must exactly match the gap between the two levels — no more, no less:

E = hf

where E is the photon's energy, h is Planck's constant, and f is the light's frequency. This is the same relationship used in the photoelectric effect, now describing an atomic transition rather than electron ejection.

Emission: Returning to a Lower Energy Level

When an electron falls back to a lower energy level, it releases a photon carrying energy equal to that gap. This is emission. The wavelength of the emitted photon determines its color in the visible spectrum, or its position elsewhere in the electromagnetic spectrum.

	Absorption	Emission
Energy	Absorbed (energy in)	Released (energy out)
Electron movement	Lower level → higher level	Higher level → lower level
Resulting state	Excited state	Ground state (or lower excited state)

Absorption vs. Emission Spectroscopy

These absorption/emission principles form the basis of two complementary analytical techniques.

Absorption spectroscopy measures the wavelengths of light absorbed as electrons jump to higher energy states. Different substances absorb light at characteristic wavelengths, which makes absorption spectroscopy useful for identifying them:

• Infrared (IR) spectroscopy studies molecular vibrations — different chemical bonds absorb specific IR wavelengths.

• UV-Vis spectroscopy examines the absorption of ultraviolet and visible light, revealing electronic transitions within molecules or atoms.

Emission spectroscopy focuses on the light emitted as excited electrons fall back to lower energy states. A familiar example is the emission of light from heated elements — such as the bright spectral lines of hydrogen seen in flame tests. These lines correspond to hydrogen's discrete, quantized energy levels, reinforcing that electronic transitions happen in fixed steps rather than continuously.

Fluorescence

Fluorescence is a related phenomenon in which a material absorbs high-energy light (such as ultraviolet radiation) and then re-emits it as lower-energy visible light. The absorbed photon excites an electron to a higher energy state, but as the electron relaxes back down, some of that energy is lost as heat rather than light. As a result, the emitted photon carries less energy — and therefore has a longer wavelength — than the absorbed photon. This shift toward longer wavelength between absorption and emission is known as the Stokes shift.

Common MCAT Mistakes

• Mixing up which direction is absorption and which is emission. Absorption always moves an electron up to a higher energy level (energy in); emission always moves an electron down to a lower level (energy out).

• Thinking any photon energy can trigger a transition. The photon's energy must exactly match the gap between two quantized levels — too much or too little energy means no absorption occurs.

• Assuming a fluorescent material emits light at the same or a shorter wavelength than it absorbed. Because some absorbed energy is lost as heat during relaxation, the emitted photon always has less energy and a longer wavelength than the absorbed photon — the Stokes shift.

• Confusing absorption spectroscopy with emission spectroscopy. Absorption spectroscopy (IR, UV-Vis) measures wavelengths a sample takes in; emission spectroscopy measures wavelengths a sample gives off, such as the spectral lines seen in flame tests.

MCAT-Style Concept Check

Question: A dye absorbs a photon of ultraviolet light and then fluoresces, emitting a photon in the visible range. Compared to the absorbed photon, the emitted photon has:

• A) Higher energy and a shorter wavelength

• B) Higher energy and a longer wavelength

• C) Lower energy and a shorter wavelength

• D) Lower energy and a longer wavelength

Answer: D

Explanation: During fluorescence, some of the absorbed photon's energy is lost as heat as the electron relaxes back toward its ground state, so the re-emitted photon carries less energy than the absorbed one. Since photon energy and wavelength are inversely related (E = hf = hc/λ), less energy means a longer wavelength — this is the Stokes shift.

FAQ

What does the Bohr model say about electron energy levels?

The Bohr model treats electrons as occupying specific, quantized orbits around the nucleus. Electrons cannot exist between these levels — they only move from one to another by absorbing or emitting a photon whose energy exactly matches the gap.

What is the difference between absorption and emission?

Absorption occurs when an electron absorbs a photon and jumps to a higher energy level (excited state). Emission occurs when an electron falls to a lower energy level and releases a photon carrying the energy difference. Both follow E = hf.

What is absorption spectroscopy used for?

Absorption spectroscopy measures the specific wavelengths of light a substance absorbs as its electrons jump to higher energy states. Because different substances absorb characteristic wavelengths, techniques like IR spectroscopy (molecular vibrations) and UV-Vis spectroscopy (electronic transitions) use this to identify and study compounds.

Why is the light emitted during fluorescence a different color than the light absorbed?

Because some of the absorbed photon's energy is lost as heat while the electron relaxes back down, the emitted photon carries less energy than the absorbed one. Lower energy means a longer wavelength, so the emitted light appears shifted toward the red end of the spectrum relative to the absorbed light — the Stokes shift.