As we go from right to left across a period, the nuclear charge increases sequentially, while atomic radius decreases as electrons in the same shell , are are held closer to the nucleus. It should therefore be harder to remove an electron from an atom further to the right on the period; and indeed it is mind you, as a physical scientist, you should look up a table of the ionization energies.
As we descend a row a column on the Periodic table , the outermost electrons are further removed from the nucleus. Ionization energies should therefore decrease down a row of the Periodic table. Does the wiki graph source illustrate what was earlier proposed?
This graph should mirror the decrease in atomic radii. Two properties are important in determining ionization energies: i nuclear charge; and ii shielding by other electrons. In partially filled shells, electrons shield each other from the nuclear charge very imperfectly, so across the Period from left to right as the nuclear charge, Z , increases, ionization energies markedly increase. On the other hand, down a Group, the increased nuclear charge is effectively shielded by the filled electronic shells.
As a chemist, as a physical scientist, you should look at some tables of ionization energies and atomic radii, and see if what I have said is reasonable. And we can see as we go down here, the number decreases. So sodium would be Potassium would be So there's a clear trend. As we go down a group in the periodic table, there is a definite decrease in the ionization energy.
So it must be easier to pull an electron away. So let's see if we can figure out the reason why. And we're going to study in detail here these two elements. So hydrogen and lithium. So let's go ahead and look at these diagrams here. We're going to fill them in for hydrogen and lithium. And so for our first diagram, we will put hydrogen.
So hydrogen has an atomic number of one. So there's one proton in the nucleus. So a plus 1 charge in the nucleus. And in a neutral atom, there's one electron. So we can go ahead and draw in hydrogen's one electron right here, like that. The electron configuration would be 1s1.
So that one electron is in an s orbital in the first energy level. So this negatively charged electron feels an attraction for this positively charged nucleus.
And so to pull it away, you must add energy. So if you add 1, kilojoules per mole of energy, you can pull that electron away. And if you do that, you'd be left with just a positive one charge in the nucleus and no electrons around it. And so you no longer have a neutral atom. You have an ion.
You have H plus, because you have a positive charge of one in the nucleus and zero electrons. So H plus. So that's the concept of ionization energy here. Let's look at lithium. So down here, we'll draw lithium. Lithium has an atomic number of three, so three protons in the nucleus. And in a neutral atom, three electrons. So the electron configuration is 1s2, 2s1. So there are two electrons in the first energy level and they're in an s orbital. So I'm going to go ahead and draw those in here.
So these two electrons I just drew represent the two electrons in the first energy level. In the second energy level, there's one more electron. So I'm going to put that electron down here like that. So for lithium, if we were to take an electron away, the one that's most likely to leave would be this outermost electron here, the one in the 2s orbital.
So if you apply kilojoules per mole of energy, you can pull away that electron. And so if you did that, you'd be left with a plus 3 charge in the nucleus. And you would still have your electrons in the 1s orbital, so I'm going to go ahead and draw those in there, but you've taken away that outer electron. And so therefore, you'd have a lithium cation here. You'd have Li plus 1, because you have three positive charges in the nucleus and only two electrons now.
So 3 minus 2 gives you plus 1. The electron configuration for the lithium cation would therefore be 1s2 because we pulled away that outer electron in the 2s orbital. So this is the picture for the ionization of hydrogen and lithium. And we're going to examine some of the factors that affect the ionization energy. And so first we'll talk about nuclear charge. Elements with a low ionization energy tend to be reducing agents and form cations, which in turn combine with anions to form salts.
Moving left to right within a period or upward within a group, the first ionization energy generally increases. As the atomic radius decreases, it becomes harder to remove an electron that is closer to a more positively charged nucleus. Conversely, as one progresses down a group on the periodic table, the ionization energy will likely decrease since the valence electrons are farther away from the nucleus and experience greater shielding. They experience a weaker attraction to the positive charge of the nucleus.
Ionization energy increases from left to right in a period and decreases from top to bottom in a group. The ionization energy of an element increases as one moves across a period in the periodic table because the electrons are held tighter by the higher effective nuclear charge.
This is because additional electrons in the same shell do not substantially contribute to shielding each other from the nucleus, however an increase in atomic number corresponds to an increase in the number of protons in the nucleus.
The ionization energy of the elements increases as one moves up a given group because the electrons are held in lower-energy orbitals, closer to the nucleus and thus more tightly bound harder to remove. Based on these two principles, the easiest element to ionize is francium and the hardest to ionize is helium.
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