However, we know today that atoms cannot be represented by quite so simple a picture. Energy Levels and Excitationīohr’s model of the hydrogen atom was a great step forward in our understanding of the atom. In this way, we now know the chemical makeup of not just any star, but even galaxies of stars so distant that their light started on its way to us long before Earth had even formed. Then they can use this knowledge to identify the elements in celestial bodies. This means that each type of atom shows its own unique set of spectral lines, produced by electrons moving between its unique set of orbits.Īstronomers and physicists have worked hard to learn the lines that go with each element by studying the way atoms absorb and emit light in laboratories here on Earth. For our purposes, the key conclusion is this: each type of atom has its own unique pattern of electron orbits, and no two sets of orbits are exactly alike. However, because these other atoms ordinarily have more than one electron each, the orbits of their electrons are much more complicated, and the spectra are more complex as well. Similar pictures can be drawn for atoms other than hydrogen. The orbital changes of hydrogen electrons that give rise to some spectral lines are shown in Figure 1. When we turn off the light source, these electrons “fall” back down from larger to smaller orbits and emit photons of light-but, again, only light of those energies or wavelengths that correspond to the energy difference between permissible orbits. Suppose we have a container of hydrogen gas through which a whole series of photons is passing, allowing many electrons to move up to higher levels. By absorbing energy, the electron can move to energy levels farther from the nucleus (and even escape if enough energy is absorbed). The closer the electron is to the nucleus, the more tightly bound the electron is to the nucleus. An electron in a hydrogen atom can only exist in one of these energy levels (or states). Thus, hydrogen atoms absorb light at only certain wavelengths and produce dark lines at those wavelengths in the spectrum we see.įigure 1: Bohr Model for Hydrogen. In this simplified model of a hydrogen atom, the concentric circles shown represent permitted orbits or energy levels. All of the other photons will stream past the atoms untouched. Only photons with these exact energies can be absorbed. Other photons will have the right energies to raise electrons from the second to the fourth orbit, or from the first to the fifth orbit, and so on. When they are absorbed, the electrons on the second level will move to the third level, and a number of the photons of this wavelength and energy will be missing from the general stream of white light. Thus, as all the photons of different energies (or wavelengths or colors) stream by the hydrogen atoms, photons with this particular wavelength can be absorbed by those atoms whose electrons are orbiting on the second level. A photon of wavelength 656 nanometers has just the right energy to raise an electron in a hydrogen atom from the second to the third orbit. Suppose a beam of white light (which consists of photons of all visible wavelengths) shines through a gas of atomic hydrogen. Let’s look at the hydrogen atom from the perspective of the Bohr model. The concept of energy levels for the electron orbits in an atom leads naturally to an explanation of why atoms absorb or emit only specific energies or wavelengths of light. We can use Bohr’s model of the atom to understand how spectral lines are formed. Explain how spectral lines and ionization levels in a gas can help us determine its temperature.Describe what ions are and how they are formed. Explain how emission line spectra and absorption line spectra are formed.By the end of this section, you will be able to:
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