 So, to recap from the last video. In the 19th century, various scientists found that elements absorbed and emitted light at certain characteristic wavelengths. Each element had a distinctive pattern of wavelengths that was like a fingerprint for that element. Just so that you're quite clear about how absorption and emission spectra are collected. For a continuous or a rainbow spectrum, white light is passed through a prism or something else that's able to separate out the wavelengths and show the spectrum. That gives you that rainbow picture, because every wavelength is present. In an absorption spectrum, a gaseous sample of the element that you want to study is put between the white light source and the prism. The atoms of the element absorb some particular wavelengths, so the spectrum that you get out at the other end looks like a rainbow with black lines where the missing wavelengths were. In an emission spectrum, you take the gaseous sample of the element and you give it some extra energy somehow, usually by heating it. And then the atoms begin to emit certain specific wavelengths of light all by themselves. So the spectrum looks mostly black, but with a few bright lines in it. And the key thing is that for a given element, the black lines in the absorption spectrum and the bright lines in the emission spectrum are at exactly the same wavelengths. As you can see by looking at the two spectra here. So this was known for a long time before Rutherford discovered the nucleus, and people started thinking in earnest about how the electrons must be arranged. Rutherford pictured the electrons orbiting the nucleus like planets. But in his model, the electrons could be at any distance from the nucleus. However, in 1913, Niels Bohr came up with a new and improved picture in which he proposed that the electrons must exist in orbits that are fixed distances from the nucleus, and they can never exist in between these fixed orbits. They can only jump from one orbit to another. Now Bohr's model was entirely theoretical, and he could not explain why electrons would exist in fixed orbits, and we still can't really explain this. However, one of the beautiful aspects of this model was that suddenly there was an explanation for the atomic spectra that had been bugging people for so long. So Bohr said, the energy that an electron has is largely to do with how close it is to the nucleus. So it's lowest in energy when it is closest to the nucleus, because it's experiencing an attraction to the nucleus. To pull it further away from the attraction of the nucleus, energy must be put in. Now Bohr referred to his fixed orbits as energy levels. He said that if an electron is to move to an orbit, a level that's further away from the nucleus, it must absorb exactly the right amount of energy to do it. This exact amount of energy is known as a quantum of energy, a precisely known amount. The wrong amount of energy and the electron either wouldn't make it to the next level, or it would get stuck in the forbidden no man's land between levels, which couldn't happen, because it's forbidden. So where does the electron get the energy from, this precise amount of energy? Well, it absorbs a photon that has exactly the right wavelength to give it the exact quantum of energy that it needs to jump up a level. So what are the missing black lines in an absorption spectrum then? They are the photons that happen to have exactly the right wavelengths to make an electron jump up a level in that particular atom. Those photons were absorbed and so they never made it through to the detector, so they're missing from the spectrum. And what about electrons coming down a level that move closer to the nucleus? Well, if orbits that are closer to the nucleus are lower in energy, then the electron must lose energy as it comes down a level. And how does it rid itself of these energy? Well, it spits out, it emits a photon. And if the electron had to absorb a certain amount of energy to move up a level, then it will emit that exact same amount of energy when it moves down a level. So the photon spat out by the electron when it moves down is of exactly the same wavelength as the photon it absorbed when it moved up. This means that if you have some atoms and their electrons are in higher energy levels and are able to move down to lower energy levels, they will do so by emitting photons of certain wavelengths. So the emission spectrum of an atom contains only lines at wavelengths that correspond to the energy differences between the orbits. The inverse, in fact, of the absorption spectrum. Now, I'm perhaps making this sound as though the electron actually means to jump up a level and goes looking for the right photon to do it. Of course, it's not like that. If the electron happens to absorb a randomly passing photon that gives it the correct energy, then it may jump up a level. But it's also entirely possible that it may absorb a photon that doesn't give it the correct energy to move levels. If this happens, it simply emits the photon again, it's a useless photon, and the electron's position remains unchanged. So just to sum up, Bohr's model of fixed electron orbits around the nucleus finally produced an explanation for the lines in atomic spectra. Absorption spectra were produced when electrons absorbed energy in the form of photons and moved up one or more levels. And emission spectra were produced when electrons that were already in higher levels emitted their excess energy in the form of photons and moved down one or more levels.