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A photon is produced whenever an electron in a higher-than-normal orbit falls back to its normal orbit. During the fall from high energy to normal energy, the electron emits a photon -- a packet of energy -- with very specific characteristics.
To "consume" a photon, meaning to absorb its energy, a molecule or atom needs to interact with the photon, causing an electron within the molecule to jump to a higher energy level, effectively absorbing the photon's energy; this process is called "photon absorption" and is fundamental to how plants use sunlight through photosynthesis or how our eyes detect light.
Key points about photon absorption:
Electron excitation:
When a photon is absorbed, its energy is transferred to an electron, which moves to a higher energy level within the atom or molecule.
Wavelength and energy:
Different wavelengths of light correspond to different photon energies, and only photons with the right energy can be absorbed by a specific molecule.
Material properties:
The ability of a material to absorb photons depends on its molecular structure and the energy levels of its electrons.
Examples of photon consumption:
Photosynthesis:
Chlorophyll molecules in plants absorb photons from sunlight, using the energy to convert carbon dioxide and water into glucose.
Vision:
When light hits the retina in your eye, photons are absorbed by photoreceptor cells, triggering a signal that is sent to the brain as visual information.
Fluorescence:
Certain molecules absorb photons at one wavelength and then emit photons at a longer wavelength, creating a glowing effect.
Absorption of Light by Hydrogen
A hydrogen atom is very simple. It consists of a single proton in the nucleus, and one electron orbiting the nucleus. When a hydrogen atom is just sitting around without much energy, its electron is at the lowest energy level. When the atom absorbs light, the electron jumps to a higher energy level (an “excited state”). It can jump one level or a few levels depending on how much energy it absorbs.
The interesting thing is that the electron can move only from one energy level to another. It can’t go partway between levels. In addition, it takes a very discrete amount of energy—no more, no less—to move the electron from one particular level to another.
The energy that an electron needs in order to jump up to a certain level corresponds to the wavelength of light that it absorbs. Said in another way, electrons absorb only the photons that give them exactly the right energy they need to jump levels. (Remember when we said that photons only carry very specific amounts of energy, and that their energy corresponds to their wavelength?)
The babble here makes it seem as if
1. Electrons ever stop moving
2. The energy defined as a color of light which is 3. Color is defined as another electron frequency is the definition of energy required to move the same variable electric charge to one of four possible states relative to the proton at center.
The absorption spectrum of hydrogen shows the results of this interaction. In the visible part of the spectrum, hydrogen absorbs light with wavelengths of 410 nm (violet), 434 nm (blue), 486 nm (blue-green), and 656 nm (red). Each of the absorption lines corresponds to a specific electron jump. The shortest wavelength/highest energy light (violet 410 nm) causes the electron to jump up four levels, while the longest wavelength/lowest energy light (red 656 nm) causes a jump of only one level.
Four-part infographic illustrating the relationship between the wavelength (color) of light absorbed by an electron, the change in energy level of the electron, and the absorption lines in a spectrum. The graphic includes: (1) a simple diagram of a hydrogen atom; (2) a graph showing the relationship between color absorbed and electron jumps; (3) an illustration of the hydrogen absorption spectrum; and (4) a graph of hydrogen’s absorption spectrum. Click View Description for more details.
View Description
The relationship between a hydrogen atom and its absorption spectrum. (Left) A simple model of a hydrogen atom showing four of the many possible “jumps” the electron could make when it absorbs light. (Right) The relationship between the electron jumps and the specific wavelengths of light that the atom absorbs. An electron jumps from one energy level to another only when it absorbs a very specific wavelength of light (i.e., when it absorbs a photon with a specific energy). The shorter the wavelength, the higher the energy, and the higher the jump. This illustration shows a set of jumps that correspond to absorption of visible wavelengths (the Balmer Series). Get the full hydrogen absorption infographic. Credit: NASA, ESA, and L. Hustak (STScI).
Emission of Light by Hydrogen
Electrons can also lose energy and drop down to lower energy levels. When an electron drops down between levels, it emits photons with the same amount of energy—the same wavelength—that it would need to absorb in order to move up between those same levels. This is why hydrogen’s emission spectrum is the inverse of its absorption spectrum, with emission lines at 410 nm (violet), 434 nm (blue), 486 nm (blue-green), and 656 nm (red). The highest energy and shortest wavelength light is given off by the electrons that fall the farthest.
The relationship between a hydrogen atom and its emission spectrum. (Left) A simple model of a hydrogen atom showing four of the many possible “drops” the electron could make when it emits light. (Right) The relationship between the electron drop and the specific wavelengths of light that the atom emits. When an electron drops down from one energy level to another, it emits a very specific wavelength of light (i.e., it emits a photon with a specific energy). The farther the drop, the shorter the wavelength and the higher the energy of the photon. This illustration shows a set of drops that correspond to emission of visible wavelengths (the Balmer Series). Get the full emission absorption infographic. Credit: NASA, ESA, and L. Hustak (STScI).
Other Atoms and Molecules
Different elements have different spectra because they have different numbers of protons, and different numbers and arrangements of electrons. The differences in spectra reflect the differences in the amount of energy that the atoms absorb or give off when their electrons move between energy levels.
Molecules, like water, carbon dioxide, and methane, also have distinct spectra. Although it gets a bit more complicated, the basic idea is the same. Molecules can absorb specific bands of light, corresponding to discrete changes in energy. In the case of molecules, these changes in energy can be related to electron jumps, but can also be related to rotations and vibrations of the molecules
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