This illustration shows a set of jumps that correspond to absorption of visible wavelengths (the Balmer Series). The shorter the wavelength, the higher the energy, and the higher the jump. 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). (Right) The relationship between the electron jumps and the specific wavelengths of light that the atom absorbs. (Left) A simple model of a hydrogen atom showing four of the many possible “jumps” the electron could make when it absorbs light.
![emission spectra emission spectra](https://cdn.comsol.com/wordpress/2016/01/Emission-spectrum-of-fluorescent-bulb.png)
The relationship between a hydrogen atom and its 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.
![emission spectra emission spectra](https://media.cheggcdn.com/media/dc7/dc799df1-0189-4f86-a108-2cb9cdd7be28/phpzAP9Iz.png)
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”).
![emission spectra emission spectra](https://i.ytimg.com/vi/jGqjRjcrqhI/maxresdefault.jpg)
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. Atomic emission spectra were more proof of the quantized nature of light and led to a new model of the atom based on quantum theory.Let’s go back to simple absorption and emission spectra. White light viewed through a prism and a rainbow are examples of continuous spectra. This would result in what is known as a continuous spectrum, where all wavelengths and frequencies are represented. Likewise, when the atoms relaxed back to a lower energy state, any amount of energy could be released. According to classical physics, a ground state atom would be able to absorb any amount of energy rather than only discrete amounts. (Credit: Christopher Auyeung, using emission spectra available in the public domain Source: CK-12 Foundation H spectrum: Commons Wikimedia, Emission Spectrum- H(opens in new window) visible spectrum: Commons Wikimedia, Linear Visible Spectrum (opens in new window) He spectrum: Commons Wikimedia, Helium Emission Spectrum(opens in new window) Fe spectrum: Commons Wikimedia, Emission Spectrum-Fe(opens in new window) License: CC BY-NC 3.0(opens in new window))Ĭlassical theory was unable to explain the existence of atomic emission spectra, also known as line-emission spectra.