In this video we'll look at another property of atoms that exhibits a periodic trend, the ionization energy of an atom. Imagine the process of taking a neutral atom and removing an electron from it. The reason electrons stay in an atom is because they're negatively charged and experience an electrostatic attraction with the positively charged nucleus. That means that in order to remove an electron from an atom you have to put in sufficient energy to break that attraction. This process is called ionization and it's always endothermic. The first ionization for an atom is when you take the neutral atom and you remove one electron from it, which forms a plus one ion. As you remove successive electrons you also have the second ionization and the third ionization and so on. Until you run out of electrons. The ionization energy is defined as the amount of energy needed to ionize a mole of atoms. So it's expressed in kilojoules per mole and different elements have different ionization energies. Potassium, this is potassium here, burning in air. Potassium's first ionization energy is 419 kilojoules per mole, which may seem like a lot. But the first ionization energy of gold is 890 kilojoules per mole, more than twice as much. In fact, 419 kilojoules per mole is quite a low value for an ionization energy. That has a really important effect. Because less energy is required to turn potassium into an ion, it's much easier for potassium metal to undergo chemical reactions in which it becomes an ion and forms ionic compounds. As a result, potassium is never found in nature as an element, only as part of compounds like potassium carbonate, potassium chloride and potassium hydroxide. If a neutral potassium atom were to be produced out there in the wild, it would find something to react with, like oxygen or water, pretty quickly. Now consider gold. It's one of the few metals, along with copper and lead, that were discovered in their metallic state early in human history. And this is because they can be found in the earth in their native elemental state, not disguised in a compound, but obviously shiny and metallic. And one of the reasons for this is that their first ionization energies are high, which makes it less likely that they will react with other chemicals around them. 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 effect reinforce each other. And it's worth saying that you need to look at these three factors altogether 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 aldermost energy level. Moving from left to right across a period, the valence shell doesn't change, so the second and the 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 of oxygen is slightly lower than nitrogen. I'm going to leave that as a problem for you to think about and research. You might like to go back over what you know about atomic orbitals, SPD orbitals, etc. And electron configurations as you ponder this.