In this video, we will learn how to describe and explain the ionization energy of elements and ions.
In chemistry, to make a positive ion, we remove at least one electron from a neutral atom. For example, we can remove an electron from a sodium atom to get a sodium ion with a one plus charge. But why do we remove one electron to make an Na+ ion instead of removing, say, two electrons to make an NA2+ ion? The answer has to do with ionization energy or the energy it takes to remove an electron. Some electrons, like the electron pictured here, will have a low ionization energy and will be relatively easy to remove. Other electrons, like the next electron we would remove from sodium to make a sodium two plus ion, will have very high ionization energies and will require a lot of energy to remove.
The most common type of ionization energy we talk about is the first ionization energy, which is the minimum energy required to eject an electron out of a neutral gaseous atom in its ground state. We can write the ionization of sodium as an equation. The neutral gaseous sodium atom produces a sodium ion and an electron, but this reaction does not occur spontaneously. We need to provide energy for it to proceed. In this case, the minimum amount of energy required to knock off this electron is 496 kilojoules per mole of sodium. 496 kilojoules per mole is the first ionization energy of sodium. We call this value the first ionization energy because the first electron is being removed.
The definition specifies that the first ionization energy has to do with a neutral atom. To find the second ionization energy or the energy it takes to remove the second electron, we can repeat the process again, starting with an Na+ ion. In this case, the minimum energy required to eject this electron is 4562 kilojoules per mole. This value is the second ionization energy of sodium because it involves the removal of the second electron.
Comparing the two values, we can see that sodium’s second ionization energy is much higher than its first ionization energy. Since it takes much less energy, it’s easier to remove the first electron from sodium than the second. This partially explains why we see more sodium one plus ions than sodium two plus ions. We will learn more about why sodium’s second ionization energy is so much higher later on in the video.
If we keep removing electrons and measuring the energy required, we can get the third, fourth, fifth ionization energies and so on. In the case of sodium, the third ionization energy is 6912 kilojoules per mole, another relatively high value especially when compared to the first ionization energy of sodium. When talking about ionization energy, the first ionization energy comes up quite frequently, while the second ionization energy and beyond are mentioned much less frequently. For this reason, even though they’re not exactly the same, we often use the term ionization energy to refer specifically to the first ionization energy. For example, we might say the ionization energy of sodium is 496 kilojoules per mole.
Now that we’ve learned about ionization energies, let’s take a look at the periodic table to compare the ionization energies of various elements. Different elements will have different ionization energies. We can look at the periodic table to find some trends in ionization energy. The first trend is that, in general, ionization energy decreases as we move down a family in the periodic table. So if we compare two elements in the same group, such as lithium and sodium, we can accurately predict that the ionization energy of the element further up on the periodic table, in this case lithium, will be greater than the ionization energy of the element further down on the periodic table, in this case sodium.
But why does this trend exist? The basic relationship to help us understand patterns in ionization energy is this. The stronger the attraction between the nucleus and the electron of interest, the more energy it will take to pull away that electron, and thus the higher the ionization energy will be. If we look at electron shell diagrams for lithium and sodium, we can see why sodium would have a weaker nucleus electron attraction. In both atoms, it’s the valence electron or the single outer electron that we’re taking a closer look at.
The first difference between these two atoms is that in the sodium atom, there’s a greater distance between the valence electron and the nucleus. In other words, sodium has a greater atomic radius than lithium. An electrostatic attraction between a positively charged thing and a negatively charged thing is stronger when those two things are closer together. So in this case, a greater distance means a weaker attraction. We might also notice that that extra distance isn’t just empty space. There’s an entire extra shell of electrons in between the nucleus and the valence electron. This extra ring of electrons introduces something we call a shielding effect. The presence of extra charged particles in between the positively charged nucleus and the negatively charged electron of interest dampens the attraction between those two particles.
Both of these two differences, the greater distance between the electron and nucleus and the shielding effect of the inner electrons, weaken the attraction between the nucleus and the electron of interest and therefore lower the ionization energy of sodium. For any two elements in the same family, the one lower down on the periodic table will have a greater distance between the electron and nucleus and more of a shielding effect from its inner electrons. So the pattern of ionization energy decreasing down the periodic table will hold.
Now, let’s take a look at the horizontal trend in ionization energy. In general, ionization energy increases as we move left to right, although there are a couple exceptions to this pattern that we will cover in a moment. To investigate the nature of this trend, let’s take a look at the elements cobalt and nickel. The most significant difference between cobalt and nickel when it comes to their ionization energies is that nickel has an extra proton in its nucleus. Within a period on the periodic table, more protons means a smaller atomic radius as the extra protons pull the outer electrons closer to the nucleus.
Similar to the first effect we looked at between lithium and sodium, a smaller atomic radius means a stronger attraction between the nucleus and any valence electron, as there’s less distance separating the two particles. A stronger attraction between the nucleus and electron means it will take more energy to remove that electron. So the element has a higher ionization energy. In this case, nickel will have a higher ionization energy than cobalt. And most elements will have a higher ionization energy than elements to their left on the periodic table. However, as we mentioned before, there are a few exceptions to this pattern.
We can graph the first ionization energy of period two elements in order to visualize the exceptions to the trend. While, in general, ionization energy increases to the right in a period, the first exception to that trend, we might notice, is that beryllium has a higher ionization energy than boron. This pattern will hold from period to period. In any period, the ionization energy of the group two element will be higher than that of the group three element. Why is this the case? To understand this exception, we need to take a look at the electronic configurations of the two elements. Beryllium’s valence electrons completely fill a 2s orbital. Having a full orbital is a stable, low-energy configuration.
Conversely, boron has an extra electron that goes into a 2p orbital. Having full orbitals adds stability. So having an extra electron that doesn’t fill the orbital is a less stable, higher-energy arrangement. It takes more energy to remove an electron from a stable configuration than from an unstable configuration. So beryllium has a higher ionization energy than boron. If we visualized the electron configuration with a diagram, we can see even more clearly how boron’s fifth electron only partially fills the 2p orbital. The other exception to the trend that we might notice is that nitrogen has a higher ionization energy than oxygen. This exception also holds in other periods. Within any period, the group five element will have a higher ionization energy than the group six element.
It’s worth noting that when we include the transition metals, we sometimes refer to these groups as group 15 and group 16 instead of group five and six. Whatever number we give them, we’re referring to the groups that contain nitrogen and oxygen. To understand this exception, we again need to look at the electron configurations of the two elements. Nitrogen’s seven electrons are distributed as follows: two in the 1s orbital, two in the 2s orbital, and the remaining three spread out among the three orbitals in the 2p subshell. Having a half-full subshell like this is a stable, low-energy arrangement.
Oxygen has a similar electron arrangement, with one key difference. Its eighth electron doubly fills one of the 2p orbitals. When two electrons occupy the same orbital, there is a repulsive force between them as they each have negative charges in close proximity to one another. Because of this repulsive force, the electron arrangement of oxygen is less stable with higher energy. It becomes easier to remove this extra electron, and thus oxygen has a lower ionization energy than nitrogen. And as an extension, the instability of this extra electron in a p subshell in any group six element lowers its ionization energy below what would be expected from the trend. In general, a group six element will have a lower ionization energy than the group five element just preceding it on the periodic table.
If we made similar graphs that show the ionization energy of the elements in different periods, they would have a similar shape, with decreases in ionization energy from the second element to the third element and from the fifth element to the sixth. Earlier in the video, we learned that sodium’s first ionization energy is relatively much smaller than its second and third ionization energies. We also learned that the arrangement of electrons within orbitals can dictate high and low ionization energies. In this case, it makes sense that sodium has a low first ionization energy because it has one more electron than neon, a noble gas with an extremely stable electron configuration of full orbitals.
If we remove one electron from sodium to create a sodium ion, we have given our sodium ion an electron configuration identical to neon’s, a very stable configuration with a full outer electron shell. Removing any more electrons would go from this stable state to an unstable state and require a lot of energy. This electron configuration explains why sodium’s second and third ionization energies are much higher than its first ionization energy. Based on this pattern, what would we expect to see in the first, second, and third ionization energies of magnesium, the next element on the periodic table? Well, compared to neon, magnesium has two extra electrons, both in the 3s orbital. We can take away these two extra electrons to form the most common ion of magnesium, an Mg2+ ion. Removing these electrons gives the ion the same stable electron configuration as neon.
Based on these electron configurations, we can make some predictions about the ionization energies of magnesium. Since the first two electrons are easily removed to create an Mg2+ ion, the first and second ionization energies should be relatively low. However, removing any additional electrons will be extremely difficult, making the third ionization energy of magnesium relatively high. If we look at the actual ionization energies for magnesium, it matches this predicted pattern. The first and second ionization energies of 738 kilojoules per mole and 1451 kilojoules per mole are relatively small when compared to the third ionization energy of 7733 kilojoules per mole.
All of the alkali earth metals in group two of the periodic table give up two electrons to form two plus ions. So another alkali earth metal, like calcium, will show the same pattern in first, second, and third ionization energies as magnesium. We see a relatively low first and second ionization energy of 590 kilojoules per mole and 1145 kilojoules per mole and then a sharp increase at the third ionization energy of 4912 kilojoules per mole. This pattern continues across the periodic table. Phosphorus has five extra electrons when compared to neon. So there’s a steady increase in its first five ionization energies, from 1012 kilojoules per mole up to 6274 kilojoules per mole. However, there’s a sharp increase from the fifth to the sixth ionization energy, up to 21268 kilojoules per mole.
The takeaway here is that there is a predictable, sharp increase in ionization energy. For group one elements like sodium, that increase happens between the first and second ionization energies. For group two elements like magnesium and calcium, it happens when we reach the third ionization energy. And we can extend this pattern across the periodic table. A group five or group 15 element like phosphorus will see a sharp increase at its sixth ionization energy. This pattern is so consistent that we can look at a list of successive ionization energies for an unknown element and use the location of the sharp increase to predict what group that element belongs to.
Now that we’ve learned about ionization energy, let’s review the key points. Ionization energy is the energy required to eject or remove an electron, specifically from a neutral gaseous atom in its ground state. The first, second, or third ionization energies refer to the removal of the first, second, or third electron. Ionization energy decreases as we move down a group in the periodic table due to those elements having a larger atomic radius and a greater shielding effect from the inner shells of electrons.
Ionization energy also increases across a period due to the smaller atomic radius that emerges when we add more protons to the nucleus. There are a couple exceptions to these patterns. Elements in groups three and six on the periodic table have lower ionization energies than predicted due to the fact that their electrons partially fill orbitals in an unstable way. Lastly, groups show predictable increases in successive ionization energies based on the number of valence electrons.