# Electron affinity and electronegativity relationship advice

### Ionization Energy and Electron Affinity

Electronegativity vs Electron Affinity The transfer of one electron from one atom to another is a very common occurrence that we do not notice. Abstract: A simple relation is presented to calculate orbital electronegativity, viz., x ~=,xo~ (1 + o. One possible definition of electronegativity is related to the mathematical average of ionization energy and electron affinity. 1 Recommendation Electronegativity is by definition the pull an atom has on the electrons in a. Electron affinity is defined as the energy released when an electron is . and relationships of the trends in atomic size, ionization energy.

Two trends are apparent from these data.

### Electronegativity (video) | Khan Academy

In general, the first ionization energy increases as we go from left to right across a row of the periodic table. The first ionization energy decreases as we go down a column of the periodic table. The first trend isn't surprising.

We might expect the first ionization energy to become larger as we go across a row of the periodic table because the force of attraction between the nucleus and an electron becomes larger as the number of protons in the nucleus of the atom becomes larger.

The second trend results from the fact that the principal quantum number of the orbital holding the outermost electron becomes larger as we go down a column of the periodic table. Although the number of protons in the nucleus also becomes larger, the electrons in smaller shells and subshells tend to screen the outermost electron from some of the force of attraction of the nucleus.

## What is the difference between electronegativity and electron affinity?

Furthermore, the electron being removed when the first ionization energy is measured spends less of its time near the nucleus of the atom, and it therefore takes less energy to remove this electron from the atom. Exceptions to the General Pattern of First Ionization Energies The figure below shows the first ionization energies for elements in the second row of the periodic table.

Although there is a general trend toward an increase in the first ionization energy as we go from left to right across this row, there are two minor inversions in this pattern.

The first ionization energy of boron is smaller than beryllium, and the first ionization energy of oxygen is smaller than nitrogen. These observations can be explained by looking at the electron configurations of these elements. The electron removed when a beryllium atom is ionized comes from the 2s orbital, but a 2p electron is removed when boron is ionized.

This, this electron can be shared in conjunction with this electron for this hydrogen. So that hydrogen can kind of feel like it's using both and it gets more stable, it stabilizes the outer shell, or it stabilizes the hydrogen. And likewise, that electron could be, can be shared with the hydrogen, and the hydrogen can kind of feel more like helium. And then this oxygen can feel like it's a quid pro quo, it's getting something in exchange for something else.

It's getting the electron, an electron, it's sharing an electron from each of these hydrogens, and so it can feel like it's, that it stabilizes it, similar to a, similar to a neon. But when you have these covalent bonds, only in the case where they are equally electronegative would you have a case where maybe they're sharing, and even there what happens in the rest of the molecule might matter, but when you have something like this, where you have oxygen and hydrogen, they don't have the same electronegativity.

Oxygen likes to hog electrons more than hydrogen does. And so these electrons are not gonna spend an even amount of time. Here I did it kind of just drawing these, you know, these valence electrons as these dots.

But as we know, the electrons are in this kind of blur around, around the, around the actual nuclei, around the atoms that make up the atoms. And so, in this type of a covalent bond, the electrons, the two electrons that this bond represents, are going to spend more time around the oxygen then they are going to spend around the hydrogen. And these, these two electrons are gonna spend more time around the oxygen, then are going to spend around the hydrogen.

And we know that because oxygen is more electronegative, and we'll talk about the trends in a second. This is a really important idea in chemistry, and especially later on as you study organic chemistry.

Because, because we know that oxygen is more electronegative, and the electrons spend more time around oxygen then around hydrogen, it creates a partial negative charge on this side, and partial positive charges on this side right over here, which is why water has many of the properties that it does, and we go into much more in depth in that in other videos. And also when you study organic chemistry, a lot of the likely reactions that are going to happen can be predicted, or a lot of the likely molecules that form can be predicted based on elecronegativity.

And especially when you start going into oxidation numbers and things like that, electronegativity will tell you a lot.

## Electron affinity: period trend

So now that we know what electronegativity is, let's think a little bit about what is, as we go through, as we start, as we go through, as we go through a period, as say as we start in group one, and we go to group, and as we go all the way all the way to, let's say the halogens, all the way up to the yellow column right over here, what do you think is going to be the trend for electronegativity?

And once again, one way to think about it is to think about the extremes.

Periodic Trends: Electronegativity and Electron Affinity

Think about sodium, and think about chlorine, and I encourage you to pause the video and think about that. Assuming you've had a go at it, and it's in some ways the same idea, or it's a similar idea as ionization energy. Something like sodium has only one electron in it's outer most shell. It'd be hard for it to complete that shell, and so to get to a stable state it's much easier for it to give away that one electron that it has, so it can get to a stable configuration like neon.

So this one really wants to give away an electron. And we saw in the video on ionization energy, that's why this has a low ionization energy, it doesn't take much energy, in a gaseous state, to remove an electron from sodium. But chlorine is the opposite. It's only one away from completing it's shell. The last thing it wants to do is give away electron, it wants an electron really, really, really, really badly so it can get to a configuration of argon, so it can complete it's third shell.

So the logic here is that sodium wouldn't mind giving away an electron, while chlorine really would love an electron. So chlorine is more likely to hog electrons, while sodium is very unlikely to hog electrons.

So this trend right here, when you go from the left to the right, your electronegativity, let me write this, your getting more electronegative. More electro, electronegative, as you, as you go to the right. Now what do you think the trend is going to be as you go down, as you go down in a group? What do you think the trend is going to be as you go down?

Well I'll give you a hint. Think about, think about atomic radii, and given that, pause the video and think about what do you think the trend is? Are we gonna get more or less electronegative as we move down? So once again I'm assuming you've given a go at it, so as we know, from the video on atomic radii, our atom is getting larger, and larger, and larger, as we add more and more and more shells.