Now that we've seen the interconnection between photons and electrons, and we've seen how they can interchange their energy, we're going to look at an emission spectrum of a hydrogen atom. And we're going to see how this emission spectrum led to our first understanding of the electronic structure in an atom. Let's begin, with what a hydrogen spectrum is. So what is an emission spectrum? This is a continuous or a line spectra of radiation emitted by a substance, continuous would be like a rainbow seeing all the various colors laid out. A line spectra just gives you individual lines at various wavelengths. Now the way you obtained this emission spectra is to energize the sample until it produces light. So you might heat it, you might pass electrical current through it, but eventually it will glow, it'll produce light, that light that's produced is passed through a prism. That rainbow that is produced, is the spectrum, and you'll have some way of visually seeing that, or electronically seeing that, it is not necessarily in the visible region so I have rainbow in quotes, it could be outside, it could be in the infrared region, it could in the ultraviolet region, it could be in the x-ray region for that matter, so wherever it is you'll see the different wavelengths separated out from each other. Here is a schematic of doing it for hydrogen, so you've got the hydrogen gas inside a glass tube on the left hand side of your screen, that tube is then got electrical current passing through it, and that electrical current then is causes it to glow, and once it's glowing, the light is passed through a slit, and focused in on a prism, and then we see the various rainbow separation that's occurring here, so we have, various slides, now this this is the visible range, and there are lines outside of the visible region, and so you could have some film that would develop at those ranges and you would see the different wavelengths of light being produced. Now here is a helium spectrum, which is different from the hydrogen spectrum, it's got a few more lines, the barium spectrum has even more lines. Now if you pass white light through the prism, you'd see all the various colors of the rainbow that you see there at the bottom, so we're just producing light at very specific wavelengths. So how did Bohr take that information and start developing his, our first understanding, of the atom. Well what he found is if he took, a wavelength there, you could take that wavelength and you could figure out the energy of that photon. And you could do that for each spot, so we could get the energy of the photon here, we could get the energy of the photon here. And we know that it's not, every possible wavelength producers, only very specific ones are produced, and there's no way to change where that light shows up on the spectrum. So, Bohr took that information and started thinking about electrons, and he thought that this could be explained if electrons can only have, very specific energy levels, and this will be re, represented by these orbits that we see going around our nucleus, so here is our nucleus of an hydrogen atom, and we've got our positive proton there. And n equals 1, n equals 2, n equals 3 here represent where the electron could be and it couldn't be anywhere in between there, so it's not existing here but it could be at 1, at 2, at 3, at 4, and at 5. Now the electron would be in its lowest energy if it could be down here in n equals 1, and that's where it wants to be. When electron is excited to an outer, spot, let's say it's excited out here to the n equals 5, it can relax back down to the n equal 1 level, now, it can relax back down to any level closer in, so it can go here, an electron from here could go there, an electron could here go all the way down the n equals 1 level. But, every transition will release one photon of light for the energy needed to make that transition, so that's key there. So, for this photon of light, and the energy associated with that photon of light, it is exactly the right amount of energy for an electron to move from here, the n equals 3, down to, the n equals 2 level, that happens to be the size of that gap. So, the light that we see here has just the right amount of energy for an electron to be going from the n equals 4 down to the n equals 2, so that's from here down to the n equals 2. This one is for this one here, from the n equals 5 down to the n equals 2, and that's just three of the possibilities, there is a lot more possibilities, but they don't fall within the visible range. If it's going to go from the n equal 5, here, down to the n equal 1, and one fell swoop there all the way down to that level, it's going to be a higher energy, and it's going to be somewhere out here beyond the visible region. And if electron goes from the n equals 2 down to the n equals 1, that's going to be a lower energy, and that might be somewhere down here, off of the visible range at a lower energy. Bohr gave these concentric circles, these orbits numbers, representing principal quantum numbers, 1, 2, 3, 4, et cetera, and they keep getting closer and closer the further away you go, so the next one will be even closer, n equals 6, and then a little bit closer together, and n equals 7, a little bit closer together, n equals 8, so they're getting closer and closer together as you move further and further away, but that is his proposed model. Now, this model didn't hold up once you put another electron in, we're going to talk more about this, we're going to talk more about calculations at the energy in a later learning objective, but this is our first understanding of the electronic structure of an atom, that is being, proposed because of the line spectrum that was produced when the hydrogen gas glows.
