Principle of Atomic Absorption and Emission: The Secrets of Light

Scientific Summary

This article traces the evolution of our understanding of light, from Newton to Bohr, passing through the fundamental discoveries of spectroscopy. It establishes that light has a wave-particle duality (photons). The energy of a photon is quantified by E = hν = hc/λ: the shorter the wavelength, the higher the energy. Young's double-slit experiment confirms the wave nature, while the photoelectric effect (Einstein) proves its particle aspect. Bohr's model introduces quantized electronic orbits, where the absorption or emission of photons corresponds to energy jumps between atomic levels, thus explaining the characteristic spectral lines of each element.

How does light interact with matter at the atomic level?

Light, the Messenger of the Universe

Light, whether from the Sun or a laboratory lamp, is the universal messenger that informs us about the composition and physics of celestial and terrestrial objects. This article on Astronoo.com traces the key steps that led to our current understanding of light-matter interaction: a journey of nearly four centuries, marked by Newton's insights, Rømer's measurements, Huygens' wave theory, and the crucial experiments of Young and Fraunhofer.

The Discovery of Photons

At the heart of this story is the discovery that light is not a continuous phenomenon, but is composed of photons, "grains" of energy whose value is rigorously defined by the Planck-Einstein relation: E = h × ν. This quantization is the key to understanding the principle of atomic absorption and emission. Indeed, electrons orbiting the atomic nucleus can only occupy specific orbits (or energy levels), like steps on a staircase.

Mechanism of Absorption and Emission

When a photon with energy exactly equal to the difference between two steps (two levels) strikes the atom, it can be absorbed, propelling the electron to a higher level (excited state). Conversely, when an electron falls back to a lower step, it emits a photon whose energy corresponds to the height of the step crossed. It is this fundamental mechanism that produces the spectral lines observed by Fraunhofer in sunlight and which today serves as a veritable "barcode" for identifying the presence of chemical elements millions of kilometers away. The article emphasizes that it is these quantized transitions that make light an unparalleled probe for exploring the Universe.

Historical Overview of the Nature of Light

Light is the only information that scientists have at their disposal to understand the world around us. Over the centuries, scientists have made light "speak."

In 1670, Isaac Newton (1643-1727) looked at the white light of the Sun through a glass prism and noticed that this beam of light was decomposed. He thought that light was composed of corpuscles.

In 1676, Ole Christensen Rømer (1644-1710) determined the speed of light by observing Jupiter's satellites.

In 1690, Christian Huygens (1629-1695) stated that light consists of a series of waves propagated through the ether, an immaterial substrate that serves as a medium in a vacuum to convey light.

In 1801, Thomas Young (1773-1829) obtained an interference pattern (image opposite), showing that light is a wave, because waves can add and subtract to create interferences (dark areas interspersed with bright areas). This experiment helped to understand the behavior and nature of light.

In 1814, Joseph von Fraunhofer (1787-1826) noticed lines in the visible light of the solar spectrum. This German optician and physicist was the first to study the diffraction of light using optical gratings (Fraunhofer diffraction). At that time, the reason for the presence of these lines in the visible spectrum of light was unknown. The answer would come much later.

In 1850, Robert Wilhelm Bunsen (1811-1899) and Gustav Robert Kirchhoff (1824-1887) discovered that the spectral lines of light emitted by an incandescent body constitute a signature allowing the identification of that body. By observing the spectrum of sunlight, they recognized several chemical elements present on Earth, including cesium and rubidium.

In 1864, James Clerk Maxwell (1831-1879) synthesized electric and electromagnetic waves. He determined that light is an electromagnetic wave and that the entire electromagnetic spectrum is light. What differentiates electromagnetic waves from each other is the wavelength. The different windows of the electromagnetic spectrum are characterized by a range of wavelengths, but also by a range of frequencies.

In 1900, Max Planck solved the black-body enigma: his formula perfectly describes the light that a body would emit according to its temperature. In other words, a high temperature indicates high energy, a low temperature indicates low energy.

In 1905, Albert Einstein (1879-1955) explained the photoelectric effect: it is the photons of the incident light that tear electrons away from matter.

Photons act as quanta of energy, which Planck had already suggested, but it was Einstein who demonstrated it. These photons therefore have a certain amount of energy, and it is this energy that removes electrons from the metal. When we receive the Sun's rays on our skin, we indeed feel the energy they carry.

Light is therefore indeed composed of photons with wave-like behavior, and each of these photons corresponds to an energy. The shorter the wavelength of the photon, the more energetic it is.

In 1911, Ernest Rutherford (1871-1937) clarified the structure of the atom and gave the atomic nucleus a size of about 10^-14 meters.

In 1913, Niels Bohr (1885-1962) proposed the structure of the hydrogen atom: electrons are located on quantized orbits. The electron navigates at a certain distance, on one of the onion-like shells around the nucleus. This is the principle of light absorption and emission in an atom.

In 1801, Thomas Young (1773-1829): the famous "Young's slits" experiment consists of causing two beams of light from the same source to interfere. This experiment, done with photons, has since been performed with all particles: with electrons in the 1920s, with neutrons in the 1950s, with atoms in the 1980s, and with molecules in the 1990s.

N.B.:
Photon energy: E = hν = hc / λ.
E is the energy expressed in joules, h is Planck's constant (6.62 × 10^-34), ν is the frequency (number of electromagnetic oscillations), c is the speed of light in a vacuum, and λ is the wavelength. The energy of a photon is therefore infinitely small. In other words, the shorter the wavelength, the higher the frequency and the greater the energy.

FAQ: Understanding Atomic Absorption and Emission

What is a photon and how is its energy calculated?

A photon is a "grain" or quantum of light, both a particle and a wave. Its energy (E) is directly proportional to its frequency (ν) via Planck's constant (h): E = h × ν. Since frequency is inversely proportional to wavelength (λ) by the relation ν = c/λ (where c is the speed of light), energy can also be written as E = hc / λ. Thus, blue light (short λ) carries more energetic photons than red light (long λ). The article specifies that this energy is "infinitely small" on a human scale, but it is sufficient to interact with electrons.

Why do atoms only absorb certain colors?

Because the electrons of an atom can only occupy well-defined energy levels (quantized orbits, according to the Bohr model). To move from one level to another, the electron must absorb or release an amount of energy exactly equal to the difference between these two levels (ΔE). If an incident photon has an energy (hν) that does not match any possible ΔE, it will not be absorbed and will pass through the matter without interaction. This is why each chemical element has a unique absorption spectrum, composed of dark lines at specific wavelengths: its luminous "fingerprint."

What is the difference between an emission spectrum and an absorption spectrum?

Emission spectrum: it is produced when an excited atom (electron on a high orbit) returns to a lower energy state. The electron then releases a photon whose wavelength corresponds to the energy difference. We observe bright lines on a dark background (for example, a sodium vapor lamp).
Absorption spectrum: it is observed when white light (containing all wavelengths) passes through a cold gas. The gas atoms selectively absorb photons whose energy corresponds to their allowed transitions, creating dark lines on a continuous background. Fraunhofer's solar spectrum is a classic example of an absorption spectrum.

How does Young's double-slit experiment prove the wave nature of light?

Young's experiment (1801) sends a beam of monochromatic light through two very thin parallel slits. On a screen placed behind, one does not observe two light spots, but a series of alternating bright and dark fringes. These "interferences" are the characteristic signature of a wave phenomenon: waves from the two slits add up (bright fringes) or cancel out (dark fringes) depending on their path difference. This experiment has been reproduced with electrons, atoms, and even molecules, demonstrating that wave-particle duality is a fundamental property of matter at the quantum scale.