 So let's look at the trends in the first ionization energy of the elements and then we'll talk about the factors that influence those values. Broadly the trend is that the ionization energy increases as you move from left to right across a period and from bottom to top up a group. If you look closely you may find that the trend is not hard and fast, you can find exceptions. You can look at those exceptions in more detail later. But for now can we rationalize that overall trend? It turns out that the three factors we talked about in the context of atomic radius also govern ionization energies. Remember them? We had the amount of charge in the nucleus or in other words the atomic number, the distance the electron is from the nucleus, what electron shell is it in, and the fact that electrons in lower energy levels can shield outer electrons from the full attraction of the nucleus. The main thing we're looking at here is how each of these affects the attraction between the nucleus and the electron that we're trying to remove. First the nuclear charge. The more positively charged the nucleus is, the greater the attractive force holding the electrons in the atom. So the more energy you have to put in to remove the electrons. So that means the greater the atomic number the higher the ionization energy in general. Second is the distance from the nucleus. Electrostatic attractions get weaker as the opposite charges move further apart rather like magnets. The further the electron is from the nucleus or in other words the higher the energy level that it's in, the weaker the attraction and the easier it is to remove. So the higher the shell or level that the electron is in the lower the ionization energy. Third is the shielding effect. As you now know an electron in a higher energy level doesn't experience the full attraction of the nucleus because there's a cloud of electrons in the lower energy levels between it and the nucleus. This shielding effect weakens the attractive force between the nucleus and the electron so the higher the energy level the electron is in the lower the ionization energy. Notice that as with the atomic radius the second and third effects reinforce each other and it's worth saying that you need to look at these three factors all together when you're analyzing ionization energies you shouldn't take them in isolation. So let's see if the trends make sense using these ideas. Remember we're talking about the first ionization energy so the electron that's being removed is one in the valence shell, the outermost energy level. Moving from left to right across a period the valence shell doesn't change so the second and third factors don't come into play. But the atomic number is increasing so the main thing that happens is that the electrostatic attraction between nucleus and electron increases. This matches the general trend. The ionization energy is increasing as you go across a period. Going down a group the atomic number increases but at the same time the valence shell is getting further from the nucleus. Hydrogen's valence shell is electron level number one, lithium's is number two, sodium's is number three and so on. This means that those second and third factors now come into play. The electron is further from the nucleus and there's more shielding from other electrons in energy levels below it so the ionization energy decreases. So what about those little glitches like that the first ionization energy of boron is actually slightly lower than that of beryllium or that that of oxygen is slightly lower than nitrogen. I'm going to leave that as a problem for you to think about and research. I might like to go back over what you know about atomic orbitals, SPD orbitals, etc. and electron configurations as you ponder this.