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Ionization Energy
We know atoms contain electrons. But what happens if we give an electron a lot of energy? We can eject it from the atom. The amount of energy needed to remove an electron from a gaseous atom is called the ionization energy (). What's the result? An ion.
A single atom can also undergo multiple ionizations. How do they distinguish each ionization? We number them the first, second, third, and so on. Remember, chemists are super creative. Each additional ionization requires more energy.
Ionization Energies and the Periodic Table
On the periodic table, generally increases from left to right in the same period. What elements have the lowest ionization energies? Group 1, the alkali metals! For them, the loss of an electron is favorable. The highest values are the noble gases. Why? As we move across a period, the nuclear charge increases. There are more protons pulling on the outer electron shell. What's the result? The atomic radius is smaller. So, it's more difficult to eject an electron.
But, there are several exceptions. Why? The ionization energy depends on the electron configuration. Atoms with a complete octet (noble gases) or filled atomic orbitals (, , ) are more challenging. Why? Those electrons are already in a stable configuration. One exception is moving from the alkaline-earth metals to the group 3 elements. For example, magnesium requires more energy than boron. Why? We need to consider the electron configurations. Magnesium has a filled s-orbital. Boron has one p-orbital electron. We know that the p-orbital is at higher energy than the s-orbital. That means less stability. Boron's electron configuration is less stable than magnesium. So, it's easier to remove its electron. Another exception occurs between groups 5 and 6. Oxygen has lower ionization energy than nitrogen. Why? Oxygen has an electron pair. At first, that may seem advantageous based on electron spin. But, there is more electron repulsion. It's a less stable configuration. So, it's easier to remove the electron.
Another trend is that values decrease going down a group. Why? We need to consider distance and shielding. Core electrons are electrons that are not in the valence shell. They shield the valence shell from the nucleus. But, electrons within the same orbitals do not protect each other. Electrons with more shielding are easier to remove. Consider the elements in group 1. They all have one outer electron in the s-orbital. But, the core electrons shield that one electron from the nucleus. The larger atoms will hold onto the outer electrons less. This occurs because there are more electrons in the atom. There is more shielding. So, it is easier to remove one. Also, the electron is further from the nucleus. We know that the electrostatic force varies with distance. So, an electron that's far away is easier to remove.
There is also a trend in multiple ionization energies. That's the removal of more than one electron from the same atom. As we remove electrons, ionization energy increases? Why? Consider the electron configuration of ions. We are trying to remove electrons from a positive ion. As it loses electrons, the nucleus holds tighter to the electrons it still has. But there's more! Let's say we removed the last p-orbital electron. Now, we're trying to remove an s-orbital electron. But, the s-orbital was already at lower energy. It was already more stable. So, it's going to be way harder to remove. As an example, let's look at carbon's ionization energies.
What do we see? There is a significant jump from the second to the third. Why? The second ionization energy extracted the last p-orbital electron. So, the third removed the first s-orbital electron. It takes a lot more energy as we go down in size. What about the massive jump between the fourth and fifth? The fourth ionization energy removes the last 2s electron. So, the fifth is the first 1s electron. It's crazy hard to remove!
