The way you learned at G.C.S.E how many electrons were present in the layers of the atom is the way that the electrons are distributed when the atom is in its unexcited state. Whenever atoms are heated or exposed to electromagnetic radiation such as light or X-rays, they become excited as they take in the energy, and that is when electrons start hopping into each others' layers like fleas.
The orbitals are always a fixed distance away from one another, and a fixed distance away from the nucleus, but the distance between each layer and the next decreases the further from the nucleus they are. In other words, the innermost layer is a certain distance from the nucleus, the next layer out is slightly nearer to the first, the next layer slightly nearer to the second etc.
The orbitals (sometimes called "layers"), can be shown as a series of straight lines showing their distance from the nucleus. You can see that they get closer and closer towards the top of the diagram. The orbitals themselves are circular, but these straight lines are a convenient way of showing them, and showing the electrons moving from one orbital to another. Obviously, it takes more energy for an electron to jump between two orbitals which are spaced far apart (such as the first and the second, or even the first and the third, since there is nothing to stop a particularly energetic electron missing one orbital out is a very high energy leap), than it does for the electron to jump between two closely spaced orbitals, say the sixth and the seventh.
Most of those outer orbitals are unoccupied. The atom only has a certain number of electrons and normally they are packed into the lowest orbitals, just as you have learned. However, just because there aren't enough electrons to fill them, it doesn't mean that the outer orbitals don't exist - it just means that they are empty.
What typically happens is this. An electron is sitting happily in its own orbital, unexcited and minding its own business, when along comes a quantum of heat energy (a small but fixed amount of energy). This energy is just the right amount (not too much, not too little) to send the electron into a higher orbital, so it absorbs the energy and jumps up. This is called energy absorption.
Electromagnetic waves can be thought of as being transmitted in miniscule bursts, called quanta (the plural of 'quantum'). These are often described as "small packets of energy", and when taken all together in their billions, gives the effect of a simple wave. A quantum of electromagnetic radiation is the smallest package of energy that you can get - they cannot be split into smaller pieces.
A quantum of energy is also sometimes called a photon. It can't really be described as a particle, nor can it really be described as a wave. It is a strange entity with both wave-like and particle-like properties. "Packet of energy" is probably the best way of describing a quantum.
When an electron has absorbed a quantum and been promoted to a higher orbital it is in an unstable state. It has left a gap in a lower orbital and its natural tendency is to fall back into that orbital. When it does this, it emits another quantum with exactly the same energy as the quantum that was absorbed. This is called energy emission.
In normal atoms, say at room temperature, electrons are continually absorbing energy as they rise from a low orbital to a higher one, and then re-emitting that energy as they fall back. The leaps that electrons take may not be matched by leaps of exactly the same size in the opposite direction. It is quite possible for the electron to make a large leap upwards skipping several orbitals as it does so, followed by two leaps downwards via intermediate orbitals:
Alternatively, it may make several small leaps upwards followed by a large leap downwards:
In this way, light of a certain energy may be absorbed and then re-emitted as light of a completely different energy. The energy of a quantum of light is proportional to its frequency (in Hz or cycles/second), with the constant of proportionality being a value called Planck's constant:
E = f.h
This constant, written as h, which was named after the German physicist Max Planck, has the value 6.62606891 x 10-34 Joules/Hz (often quoted as "Joules seconds", which, if you think about it, is actually the same thing). This simple relationship between frequency and energy leads us to believe that low frequency radiation, such as radio waves or microwaves, has little energy, whereas high frequency radiation such as X-rays and gamma rays contains a great deal of energy. This explains why generally high frequency radation can penetrate objects that low frequency radiation cannot (e.g. X-rays can penetrate human flesh whereas light rays cannot).
The electromagnetic radiation which atoms emit can only have certain energy values depending on the exact orbitals between which the electron is jumping. This means that the quanta emitted can only have certain frequencies. It is the same when electrons are absorbing energy. They can only absorb quanta of certain energies (certain frequencies), which correspond exactly to the energy needed to make a leap - they cannot absorb energy which would take them up a few orbitals "with a little left over", atoms simply don't work like that. Only quanta of exactly the right frequencies can be absorbed.
Different atoms have different numbers of electrons. These electrons absorb and emit energy in such a way that is unique for each single type of atom. A helium atom, for example, abosrbs and emits a certain pattern of energy frequencies. A neon atom will absorb and emit an entirely different pattern. A uranium atom will absorb a totally different pattern of energies still. These patterns identify the atoms from which they originate uniquely - they are like fingerprints in human beings.
The frequencies are focussed by the telescope and then strike a triangular prism. Prisms cause beams of light to change direction (the so-called "bending" of light), technically known as refraction, but only when the light strikes the surface of the prism and when it leaves the prism at the opposite side. While the light is travelling through the prism it travels in a straight line, just as it does in air. More importantly, light of different frequencies is refracted to a different extent: High frequencies are refracted more than low frequencies. This is the physical cause of white light splitting into its rainbow colours: The "red" light is refracted less than the "violet" light, so the colours spread out.
The result is that by the time the light reaches the other focussing tube, the frequencies present in it have all separated themselves and are now spread out in a line, with the highest frequencies to the left and the lowest frequencies to the right. The higher the frequency of any particular component, the further to the left it has been bent.
These frequencies are focussed separately (not recombined into one beam, just "sharpening" the individual frequencies) by the second tube just as they were by the first tube, except that now they all focus to different point, being slightly separated from each other. What you see when you peer in at the other end is a series of bright and faint lines, each with a different colour, like one of those supermarket bar codes. Each line is one frequency which has been moved to a certain point along the scale from left to right by passing through the prism and bending just so far. The reason that there are so many lines is because there are many possible ways for the electrons to jump between orbitals. Indeed, there are many lines that can't be seen at each end of the visible spectrum, as they are in the infrared or ultraviolet parts of the spectrum or even further "along the dial".
To view the absorption spectrum of an element, the atoms of the element are not heated directly. Instead a chamber containing the atoms is illuminated by a broad band of frequencies (i.e. by a complete spectrum containing all the frequencies). The electrons of the target atoms will absorb some of the bombarding frequencies, the very same frequencies that are found in the emission spectrum. Those electrons then emit the frequencies as they fall back into lower orbitals.
However, the quanta of energy are re-emitted in all directions. The result is that the total amount of energy that is passing in the original direction has been reduced - the atoms have scattered the special frequencies. When you look at an absorption spectrum, what you see is a complete rainbow spectrum with a few dark lines in it. These dark lines correspond exactly to the bright lines that you would get in an emission spectrum. In fact, absorption spectra have been described as being "photographic negatives" of emission spectra.
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