[MUSIC] The key to doing analysis at such tiny levels is spectroscopy, and spectroscopy relies on the fact that when light is passed through a suitable material, it's broken up into its component colours, which is the spectrum. And this is something we're all familiar with, we've all seen a rainbow, and this is where white light from the sun is broken up into its component colours by raindrops in the air. The same phenomenon can be achieved using a glass prism. So, here on the left you can see white light passed into a prism and broken up to give the spectrum. This is a phenomenon that was investigated centuries ago. The British mathematician and physicist Isaac Newton for instance, extensively experimented with this phenomenon. Now in those days of course, the only source of light bright enough to do these experiments was sunlight. So people like Newton would use sunlight and then observe the resulting spectrum. So as instrumentation improved, something very strange was noticed in the spectrum that was obtained from sunlight. Looking very carefully at this spectra, Fraunhofer, a German physicist, noticed that the spectrum was not continuous. Within the spectrum obtained from the sun, there were discrete black lines, and people wondered why these lines existed. And it was realized that these lines were due to elements in the atmosphere of the sun, particularly, hydrogen. So as the light came out of the sun, light at particular frequencies was being absorbed by hydrogen, and therefore, these black lines were appearing at particular points in the spectrum. Now interestingly, not all of these lines could be matched to known elements like hydrogen. Some of the lines could not be matched to any element then existing on Earth, and this led the astronomer Norman Lockyer to propose that there was an as yet unknown element present in the atmosphere of the sun - an absorption of light by that element was causing these unknown lines. So he named that element Helium, and helium became the first element discovered outside the Earth, rather than on the Earth itself. And he was proved right, because later, helium was discovered on Earth in minerals and it was found to absorb light in exactly the same way as did helium in the sun. Now, we are not restricted just to visible light. Light makes up just a very small portion of what we refer to as the electromagnetic spectrum. This is in fact a very wide ranging spectrum which goes from the low energy radio waves through microwaves into the infrared, then you have the very narrow band of the visible spectrum, and then out into the ultraviolet, and on into x-rays and gamma rays. So while we are all familiar with rainbows and the visible spectrum of light, we can actually use almost any part of the whole electromagnetic spectrum for analytical purposes. But to understand why it's possible, we have to start looking at the structure of the atom. So at the end of the 19th century, there was a model of the atom which is described as the plum pudding model. And in this model, the atom was a large positively charged object with negatively charged electrons embedded in it, rather like the raisins in a Christmas pudding. This model was dropped based on experimental work, and Rutherford proposed a new model. In his new model, the atom consisted of a small, very dense, positively charged core called the nucleus, which was later found to be made up of protons and neutrons, and almost all the mass of the atom is concentrated in that nucleus. The electrons, which are negatively charged, orbit around that nucleus. So most of the atom is free space and the electrons are moving through that free space as they orbit the nucleus. Now, there's a problem with this model which needs to be explained. If you have this small, dense, positively charged nucleus in the centre surrounded by these negatively charged electrons, then the atom should destroy itself because the positive nucleus should attract the negative electrons, and the electrons should just collapse into the nucleus and that's the end of your atom. [SOUND] But atoms don't destroy themselves; they are stable. So a new theory was needed, and this was provided by the Danish physicist, Niels Bohr, using the Quantum Theory that had originally come from the German physicist, Max Planck. And in Bohr's idea, the electrons were not free to move anywhere in the atom. They were restricted to particular levels. This is because energy is quantized. Energy is not a continuous function, but it exists in discrete little packages. Now in our everyday lives, we don't really notice this. This is because these little packages or quanta of energy are so small compared to the scale on which we live our lives that we don't notice it. But down on the atomic or molecular scale, this becomes very, very important indeed. So, Bohr said that electrons are restricted to specific energy levels. However, the electrons can move from one level to another. For instance, if the electron is provided with sufficient energy, then it can jump up to a higher energy level by absorbing that energy. Similarly, if an electron is in a higher energy level, it can drop down to a lower energy level and release that energy. The energy that's absorbed or released is often in the form of electromagnetic radiation, and sometimes in the form of light itself. Now, the amount of energy that's absorbed or released must match exactly the difference between the two energy levels. And according to the de Broglie equation, that energy is linked to the frequency of the electromagnetic radiation by a very simple formula. And if we're talking about visible light, then of course, the colour of that visible light is dictated by the frequency. [BLANK_AUDIO] Let's look at a very simple analogy for the Bohr concept. The energy levels around the atom are similar to the rungs of a ladder. If you're climbing up a ladder, you have to stand on one of the rungs, you can't stand in the space in between. So the rungs of the ladder are like those allowed energy levels. So if you're climbing a ladder and you want to go up to the next higher rung, then energy has to be provided. So the energy is absorbed by the system as you go up to the next rung. In terms of spectroscopy, it works like this. Suppose you have a sample and you're passing a beam of radiation such as light through that sample, and you have a detector on the other side. As you scan through the different frequencies, the light is not absorbed until you get to the frequency of light that matches the energy difference between those energy levels. Then the light is absorbed and the intensity of light coming out the other side of the sample drops. And then as the frequency moves on, the energy is no longer absorbed, and the intensity of light at the detector goes back to its original value. So this point in the graph where the light is observed would of course correspond to one of those black lines observed in the solar spectrum by Fraunhofer. And this is what happens in terms of the atom. When you get to the right frequency, as you scan through the light source of frequencies, the electron can absorb the energy to be promoted to a higher energy level, and this is called the Excited State. And again, the delta E, the change in energy as it goes up to the next level, is characteristic of that particular element. If we go back to our analogy of the ladder, suppose you're at a higher rung on the ladder and then you go down, then of course, energy is released. And again, that energy matches the gap between the two energy levels. In terms of spectroscopy it works like this. Suppose you have a sample and you supply some form of energy to the sample in order to promote the electrons up to higher energy levels. Then as those electrons drop down to the lower energy levels, light is emitted. If we use a detector to measure the frequency of the light, then we find that most frequencies are not emitted at all because they don't correspond to the energy level differences. But we get emission at specific frequencies which match those energy level differences. So in atomic terms, this is what's happening. The energy we supply has promoted an electron so that the atom is in its excited state. As the electron drops down to a lower energy level, energy is given out in the form of electromagnetic radiation called light. And once again, that frequency of that radiation coming out matches the energy difference between those two levels. So, here's an example of how to obtain an emission spectrum of an element. In this case, it's hydrogen. Hydrogen is in the gas discharge tube, it's electrically energized into its excited state. As the atoms in the excited state drop back down, they emit light, and if we pass that light through a prism, we obtain a pattern. We obtain a pattern that's mostly black, but with certain coloured lines which match the energy level differences. Now as this is hydrogen, if we take this emission spectrum here and we compare it to Fraunhofer's solar spectrum, we will see that some of these coloured lines from the hydrogen emission match some of the black lines in Fraunhofer's solar spectrum which come from hydrogen absorption. So we have two complementary methods. We can do absorption spectroscopy, where we're looking at what kind of light is absorbed by a particular atom as the atoms go from ground state to excited state; or we can do emission spectroscopy, where we're exciting the atoms and we monitor what kind of light is emitted as they go back down to the ground state. Now, part of the Bohr model is that there are multiple energy levels for any particular atom. And that means we're not looking at a single emission going from one single energy level to another, but we're looking for multiple emissions or ab...
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