Why is the spectra for each element unique

You can view the atomic spectrum of each element at. Here is a look at emission colors of light produced by four different elements. Video from: Noel Pauller. This video show uses diffraction grating to show the emission spectra of several elements including hydrogen, oxygen, neon and nitrogen. Use of a tool such as a spectroscope would allow someone to determine the different wavelengths each of these elements is giving off.

The color you observe in the video is the sum total of all of the visible emissions from each element. A common lab performed in chemistry involves flame tests of different metal salt compounds. Different compounds will give off different colors of light. The color can be used to identify which elements are present in the salt.

How can atomic spectra be used to identify elements? Apr 22, Explanation: When an atom absorbs energy, its electrons jump to higher energy levels. 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. Only photons with these exact energies can be absorbed. All of the other photons will stream past the atoms untouched. Thus, hydrogen atoms absorb light at only certain wavelengths and produce dark lines at those wavelengths in the spectrum we see.

Figure 1: Bohr Model for Hydrogen. In this simplified model of a hydrogen atom, the concentric circles shown represent permitted orbits or energy levels. An electron in a hydrogen atom can only exist in one of these energy levels or states. The closer the electron is to the nucleus, the more tightly bound the electron is to the nucleus. By absorbing energy, the electron can move to energy levels farther from the nucleus and even escape if enough energy is absorbed.

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. The orbital changes of hydrogen electrons that give rise to some spectral lines are shown in Figure 1. Similar pictures can be drawn for atoms other than hydrogen. 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.

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. 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. Astronomers 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.

Then they can use this knowledge to identify the elements in celestial bodies. 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.

However, we know today that atoms cannot be represented by quite so simple a picture. For example, the concept of sharply defined electron orbits is not really correct; however, at the level of this introductory course, the notion that only certain discrete energies are allowable for an atom is very useful.

Ordinarily, an atom is in the state of lowest possible energy, its ground state. In the Bohr model of the hydrogen atom, the ground state corresponds to the electron being in the innermost orbit. The atom is then said to be in an excited state. Generally, an atom remains excited for only a very brief time.

After a short interval, typically a hundred-millionth of a second or so, it drops back spontaneously to its ground state, with the simultaneous emission of light. The atom may return to its lowest state in one jump, or it may make the transition in steps of two or more jumps, stopping at intermediate levels on the way down. With each jump, it emits a photon of the wavelength that corresponds to the energy difference between the levels at the beginning and end of that jump. An energy-level diagram for a hydrogen atom and several possible atomic transitions are shown in Figure 2 When we measure the energies involved as the atom jumps between levels, we find that the transitions to or from the ground state, called the Lyman series of lines, result in the emission or absorption of ultraviolet photons.

In fact, it was to explain this Balmer series that Bohr first suggested his model of the atom. The right hand side a of the figure shows the Bohr model with the Lyman, Balmer, and Paschen series illustrated.

As these arrows are moving away from the nucleus, they represent absorption of energy by the atom to move an electron up to each level.

As these arrows are pointing toward the nucleus, energy is released from the atom as electrons. When it does this, it loses energy. The amount of energy it loses will be equal to the difference in the energy levels it moves between. This energy is released as a photon. The energy of the photon can be worked out using the equation. For this relationship:.