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. Atoms that have absorbed specific photons from a passing beam of white light and have thus become excited generally de-excite themselves and emit that light again in a very short time.
You might wonder, then, why dark spectral lines are ever produced. Imagine a beam of white light coming toward you through some cooler gas. Some of the reemitted light is actually returned to the beam of white light you see, but this fills in the absorption lines only to a slight extent. A spectral line is like a fingerprint that can be used to identify the atoms , elements or molecules present in a star , galaxy or cloud of interstellar gas.
If we separate the incoming light from a celestial source using a prism, we will often see a spectrum of colours crossed with discrete lines.
Note that spectral lines can also occur in other regions of the electromagnetic spectrum , although we can no longer use a prism to help identify them. Emission lines are seen as coloured lines on a black background. Absorption lines are seen as black lines on a coloured background. I searched on the site and discussed with various members of the community and some of the conclusions I drew are mentioned below:.
Electron-nucleus attraction and the electron-electron repulsion. Hydrogen is a special case because it has only a single electron so there is no electron-electron repulsion. Helium has two electrons so now we have some e-e repulsion, Lithium has three electrons so there is even more e-e repulsion and so on Every atom has a different number of electrons and a different nuclear charge, so the balance of nuclear attraction and e-e repulsion is different.
So every atom has its own set of atomic orbitals that are specific to it and are different from every other atom.
The spectrum comes from transitions between orbitals, and since the orbitals are different for every atom the spectrum is different for every atom. Do all noble gases or alkaline earth metals have similar spectral lines considering the above points I mentioned? Do all noble gases or alkaline earth metals have similar spectral lines considering the above points? The question is interesting after you modified it. The basic set of reasoning you provided is the main story.
Each element has a different nuclear charge and the outermost electron s is responsible for the atomic emission spectrum. Since the nuclear charge is different, those outermost electrons experience a different potential energy. Their kinetic energy is also different from element to element. The key question is what is meant by similarity? The atomic spectra of all the elements is visualized as bright lines on a dark background.
The reason they appear as lines is just because of the instrument used to observe the atomic spectrum. There is nothing fundamental in the "line"spectrum. The atomic emission appears as lines because the slit in the monochromator is shaped like a very narrow rectangle.
This is the image of the slit. If I made a very narrow circular opening, the images will appear as bright points rather than lines. Consider that line spectrum is the conventional way of looking at the atomic spectrum. Are they similar in any way? Visually there is no similarity at all in the spectrum. Look at the spectrum. The reason we cannot see any similarity is because the dispersion of the lines is on the "wavelength" scale.
However, very bright, I mean really genius spectroscopists of the 18th and the 19th century were able to find out a patterns or call it as mathematical series.
They figured out that each set of lines forms a pattern as Sharp series, Principal series, Diffuse Series and Fundamental series and there they found similarities in patterns-in terms of mathematical series.
In short, our eyes cannot interpret those series. Another complication arises because we are just looking at the visible region.
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