A blue arrow with a relatively open wave pattern representing blue medium-wavelength light travels into the atom to strike the electron.The straight white arrow points from level 2 to level 5 to show a jump of three levels. A blueish-purple arrow with a tight wave pattern representing blueish-purple short-wavelength light travels into the atom to strike the electron.The straight white arrow points from level 2 to level 6 to show a jump of four levels. A purple arrow with a very tight wave pattern representing purple short-wavelength light travels into the atom to strike the electron.In all four scenarios, the small circle is positioned on energy level 2 to indicate the electron’s starting energy level. The change in energy level is shown with a dashed, straight white arrow. Light is represented as a wavy colored arrow. In all four cases, the electron is represented as a small circle. The distance between adjacent energy levels decreases with distance from the nucleus.įour scenarios involving absorption of light and electron jumps are shown. The circles are labeled “level 1” through “level 6” with level 1 closest to the nucleus, and level 6 farthest. Six concentric circles representing electron energy levels (or orbitals) surround the nucleus. Graph of hydrogen’s absorption spectrumĭiagram of a Hydrogen Atom (left side of the infographic)Ī diagram of a hydrogen atom shows the relationship between the color of light absorbed by an electron and its change in energy level.Īt the center is a solid circle representing hydrogen’s nucleus.Illustration of the hydrogen absorption spectrum.Graphic showing the relationship between color absorbed and electron jumps.This four-part infographic titled “Absorption of Light by Hydrogen” illustrates the relationship between the wavelength of light absorbed by an electron in a hydrogen atom, the change in energy level of the electron, a picture of the absorption lines in the hydrogen spectrum, and the graph of hydrogen’s absorption spectrum. 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. Each of the absorption lines corresponds to a specific electron jump. 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). The absorption spectrum of hydrogen shows the results of this interaction. (Remember when we said that photons only carry very specific amounts of energy, and that their energy corresponds to their wavelength?) Said in another way, electrons absorb only the photons that give them exactly the right energy they need to jump levels. The energy that an electron needs in order to jump up to a certain level corresponds to the wavelength of light that it absorbs. 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 interesting thing is that the electron can move only from one energy level to another. It can jump one level or a few levels depending on how much energy it absorbs. When the atom absorbs light, the electron jumps to a higher energy level (an “excited state”). When a hydrogen atom is just sitting around without much energy, its electron is at the lowest energy level. It consists of a single proton in the nucleus, and one electron orbiting the nucleus. Absorption of Light by HydrogenĪ hydrogen atom is very simple. Why is this? Let’s take a look at hydrogen, the most abundant element in the universe. We can do both of these because each element has its own unique spectrum.Īn element’s spectrum is like its fingerprint, its autograph, its barcode. We can use a glowing nebula’s emission spectrum to figure out what gases it is made of based on the colors it emits. We can use a star’s absorption spectrum to figure out what elements it is made of based on the colors of light it absorbs. Let’s go back to simple absorption and emission spectra.
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