Video Transcript
In this video, we will learn how to define electron affinity and describe and explain trends in electron affinity throughout the periodic table. Chemists have a tough job keeping track of how much energy there is in the system, and all the different types can be exhausting. One of the ways chemists make it easier is to break things into smaller steps. If we have an atom, we can measure the energy required to remove each electron one by one until we have a bare nucleus. We call the process of removing these electrons ionization. Here are the first and second ionizations of a helium atom.
But what about the opposite? It’s possible for atoms to gain electrons as well. So what do we call that? Here is an atom of hydrogen. Typically, we’d find elemental hydrogen in hydrogen molecules H2. However, most elements don’t form molecules the way hydrogen does, so it’s hard to compare them that way. Instead, we start with separated atoms in the gas state. When a hydrogen atom is close enough to an electron, they’ll attract each other a little and they’ll stick together. Chemical potential energy will be converted to other forms of energy like heat. The energy released in this process is known as the electron affinity of hydrogen.
We also use the word electron affinity to refer to the overall process of adding electrons. Electron affinity is defined as the energy released when an electron is added to a neutral atom in the gas state to form a negative ion, and we usually express electron affinities in kilojoules per mole of atoms. There is such a thing as a second electron affinity, where the electron is added to a one minus ion instead of a neutral atom. We’ll have a quick look at this later, but we’re going to focus more on first electron affinities.
Before we go any further, it’s very important to know the relationship between electron affinity and enthalpy or energy changes. The electron affinity of hydrogen is not the same as the enthalpy change when we add an electron to a hydrogen atom. The electron affinity is the release of the energy, that’s the amount of energy coming out, while the enthalpy change is the change in energy of the system. If the electron affinity is positive, and we can use the symbol E ea, the enthalpy change for the same process is negative. This corresponds to an exothermic process, a release of energy into the surroundings, while a negative electron affinity indicates that the enthalpy change for the process is positive. This corresponds to an endothermic process.
While this convention is frustrating, you simply need to remember the definition of electron affinity. If you do that, you should be able to figure out what the sign of the electron affinity means. There are currently 118 elements on the periodic table. This means there are potentially 118 different first electron affinities for us to measure, assuming that isotopes don’t make a difference. The electron affinity of hydrogen is about 73 kilojoules per mole. This means that if we have one mole of hydrogen atoms in the gas phase and we add an electron to each of them, we’ll convert 73 kilojoules of chemical potential energy to other forms of energy like heat. So for each mole of hydrogen atoms, we turn into H‒ ions, we’ll get 73 kilojoules of energy out to the surroundings.
What this means is that an H‒ ion is more stable than a separated hydrogen atom and electron. But what about elements that we know are already very stable, like the noble gas helium? When we try and add an electron to helium, the repulsion from the electrons is greater than the attraction from the nucleus. This means adding electron to a helium atom requires us to put energy in. It’s nearly impossible to measure this directly since He‒ isn’t stable, but we can do some calculations.
If we could add an electron to a helium atom, it would take about 48 kilojoules per mole. This means that the electron affinity of helium is about negative 48 kilojoules per mole. So we’ve seen that some elements have positive electron affinities and some of them have negative electron affinities. So for some elements, adding electron to their gaseous atoms is exothermic, and for others, it’s endothermic. This is a color map for the electron affinities of the elements. The pinker the square, the greater the electron affinity of that element. The bluer the square, the lower the electron affinity. Gray indicates electron affinities that are close to zero, and white indicates that the electron affinity has yet to be determined.
As you can see, there isn’t a consistent trend, but we do see a few patterns here and there. For instance, all the noble gases have negative electron affinities. But otherwise electron affinities generally become more positive, moving left to right and bottom to top. Generally speaking, we can see an increase in the electron affinity of the elements across a single period moving left to right. And for quite a few of the groups, we can see an increase in the electron affinity as we go from the bottom to the top.
Let’s look at the trend going down a group first. This trend is strongest in groups one, 14, 15, 16, and 17. As we go down a group on the periodic table, atoms of the elements get bigger and bigger and bigger as they have more electron shells. With similar elements as the size of the atoms increases, the attraction to an extra electron will decrease. If we try to add an electron to a small atom, it’ll be able to get closer to the nucleus, while for atoms with more electrons, the electron won’t be able to get as close.
Before we move on, we are going to look at one inconsistency in group 17. Fluorine has an unusually high electron affinity at 328 kilojoules per mole, but chlorine’s is even higher. Meanwhile, the electron affinities of bromine, iodine, and so forth continue as expected, following the downward trend. While it’s hard to know exactly why this happens, we can come up with a decent theory.
A fluorine atom has a tiny atomic radius of only 42 picometers. Chlorine’s is almost double at 79, while those of bromine and iodine increase incrementally. Here is one atom of fluorine, nine electrons surrounding a nucleus with a nine plus charge. A fluorine atom is even smaller than a hydrogen atom, which has an atomic radius of 53 picometers. We can imagine that the negative charge of those electrons in a fluorine atom will be very, very dense. This is going to reduce the electron affinity of fluorine relative to chlorine because the incoming electron is experiencing greater repulsion.
Now let’s take a closer look at the trend going left to right across a period. Let’s take a closer look at period two. As we go left to right, the atomic number increases. Since the number of electron shells is not increasing and the nuclear charge is increasing, the atomic radius goes down. In an atom of lithium, the nucleus has a charge of three plus, so we have three electrons in the electron cloud. In total, these electrons have a three minus charge. Meanwhile, a fluorine atom has a nine plus charged nucleus and nine electrons. For this trend, we’re going to ignore the noble gas neon, since we already know it’s very stable and won’t accept an extra electron.
Now let’s imagine this extra electron being introduced next to a lithium or a fluorine atom. The electron approaching a fluorine atom can get much closer to the nucleus before being repelled by the electrons. So this is a basic explanation of why, in general, we see an increase in electron affinity going left to right across a period of the periodic table.
The last thing we’re going to look at, and I said we’d go back to it, is second electron affinities. Second electron affinities are associated with this process: X‒ plus an electron forms X2‒. Again this happens in the gas phase. The interesting thing here is that we’re introducing an electron which is negatively charged to an anion. So the negative ion will naturally repel the electron, raising the energy required to add them together. Because of this, second electron affinities are always negative and therefore endothermic.
We can see this with the element oxygen. The first electron affinity of oxygen is about 141 kilojoules per mole. The enthalpy change for this process is negative; we’re dealing with an exothermic process. But the second electron affinity of oxygen is predicted to be about negative 744 kilojoules per mole. So the enthalpy change for this process is positive; we’re dealing with a vastly endothermic process. Now, what’s interesting is that O2‒ ions are formed all the time in lattices with metals. This endothermic energy cost is paid for when ions of different charges come together.
Let’s finish up with the key points. Electron affinity of an element is the energy released when an electron is added to a neutral atom of that element in the gas state to form a negative ion. Electron affinities are normally expressed in kilojoules per mole. Electron affinity is often used to refer to the process itself, adding an electron to a gaseous atom. An element’s second electron affinity is the energy associated with the process of adding an electron to an ion with one minus charge. A positive electron affinity indicates a negative enthalpy change, indicating an exothermic process. And, naturally, the reverse, a negative electron affinity indicates a positive enthalpy change and an endothermic process.
On the periodic table, we see a trend of increasing electron affinity generally moving left to right and bottom to top. Within a group moving bottom to top, we attribute this trend to a decreasing atomic size. And across a period, an increase in atomic number of the element means that the atom will have a higher nuclear charge. This contracts the electron shell and allows the incoming electron to get closer, being more attracted to the nucleus. Bear in mind that often we ignore the noble gases when discussing these trends because noble gases have atoms with full valence shells. They’re extremely stable, and adding electron will always be endothermic.
