In this lesson, we're going to introduce concepts of the atomic scale structure. And what we'll begin with is describing the Bohr Model which was introduced in the early part of the 1900s. We'll look at some modifications to that in terms of the development of the understanding of the energy levels. We'll introduce the concept of orbitals that was presented to you most likely in a course of chemistry in high school and we'll describe then the order of filling of these electron configurations. So now let's begin by talking about the Bohr Model, and we'll start as an example, the zinc atom. When we look at the types of orbits that occur in Bohr Model think of this as a planetary model, we have a variety of quantum numbers that describe the various primary orbits. And we introduce a second type of atomic structure called l, which represents when l is equal to 0, the s electrons and it turns out that there are two electrons in that s shell one that has a spin up, one that has a spin down. When we go to l = 1, what we now have is p electrons and we're describing the p electrons, there are a total of 6 of those each with spin up, spin down. So there's a total then of 6 electrons. Then we introduce the concept of the d electrons. Once again, we have five levels. And in each one of those levels as we go through the filling of these shells we're going to wind up with a total of ten electrons. When we talk about l = 3, we're now describing the f electrons and it turns out that there are a total of 14 of those type electrons. So looking at the atomic number 30, which is the atomic number for zinc. The way that structure is laid out in the original Bohr Model is to look at each one of the shells and when we add the modifications to the orbital is what we see is there is on the inner ring the 1s electrons out there is a a maximum number of two of those electrons as we move further out. We find the second orbital of the 2s electrons. So we again we have a total of 2 there. We move out to the 2p electrons and it turns out those p electrons contains a total of 6 electrons in that configuration. We go continue to go out until ultimately, we wind up partially filling some of the electrons in the d, and those are represented by the electrons in the 3rd n is equals to 3 and the d electrons in that orbit. So when we look at these orbits, we need to understand what the origin is of these configurations of spdf come from. It turns out that these were spectral line descriptions with respect to individual atoms that were isolated, and the spectra of these different atoms were determined. And so when the original chemist looked at these spectral lines, it turned out that there were different ways that they could be described. Some of the lines were sharp, and therefore were determined as s. Some of the lines were what were referred to as principal lines and those were p lines. And then we had the d, which stands for diffuse and f stands for fundamental. And what I'll show next is a schematic that represents the relative energies of the spdf electrons, and that's what we have here. So as we look at the principle quantum number n which is labeled along the x axis and the energy, we see that each one of these spd's and f have different energy levels. Now what we can do is we can come up with a scheme that will help us organize the electrons as they fill across the periodic chart. And we'll be looking then at the atomic numbers and the corresponding spd electrons that are involved and the order in which these fill. So if we use the device I have indicated on the right hand side we have an arrow and so what that indicates is that the s electrons, the one s electrons will be filled. The first one is in the case of hydrogen where we have one electron. The second is helium in which we have the s in the first principle quantum number filled. Then we go to the second level, the n equal two level. And now what we have is a total of 2 electrons and this time, it is the 1s and then the 2s, and then the 2s will fill those and then we begin to fill the p. In this case, we begin by looking at 2p, and the p's have a total of 6 electrons in them. The next, in terms of the level of filling will be the 3s electrons which again have 2, if we continue on we'll have the 3p and the 4s. Ultimately we'll begin with the 3d which represent what we will see in a few minutes. The d electrons and those electrons are associated with the transition metals, and we'll see how they're grouped together. And then again we as we continue to fill, we'll start talking once again with p electrons. And then ultimately, with the s electrons. And we continue to do this, using the arrows and following through as I have indicated at the bottom of the slide. Now it turns out that there are some exceptions to this order of filling, and we will wind up discussing those. Now if we go back and consider a material like sodium, and we look at the energy levels and the way these electrons wind up filling. When we start out with the electrons in the 1s shell, those inner electrons, and again they're two of them. We go to the 2s, and now what we have is 2 electrons in that 2s shell. We continue to fill and then we completely wind up piling the p electrons with a total of 6. Remember each one of those couples represent one electron with one spin and the other the electron in the other spin. Finally, we get to 3s which represents the top electron that is in the sodium structure, and we're going to ultimately refer to this electron as a valance electron. Now when we talk about the electrons in the spd and f series, what we're actually seeing are configurations that are given by distribution functions where rather than seeing a symmetrical cloud which is what we see in the s electrons. But when we start looking at p electrons, d electrons and f electrons, what we see is the fact that these electron orbits are highly directional. So we can then see the relationship that exists. And we'll be looking at this in the next lesson, where we're going to try to show the relationship between these different orbitals and the location of the elements on the periodic chart.