What does ionization energy mean

Electron affinities are more difficult to measure than ionization energies and are usually known to fewer significant figures. The electron affinities of the main group elements are shown in the figure below. Several patterns can be found in these data. Electron affinities generally become smaller as we go down a column of the periodic table for two reasons. First, the electron being added to the atom is placed in larger orbitals, where it spends less time near the nucleus of the atom.

Second, the number of electrons on an atom increases as we go down a column, so the force of repulsion between the electron being added and the electrons already present on a neutral atom becomes larger. Electron affinity data are complicated by the fact that the repulsion between the electron being added to the atom and the electrons already present on the atom depends on the volume of the atom. As a result, these elements have a smaller electron affinity than the elements below them in these columns as shown in the figure below.

From that point on, however, the electron affinities decrease as we continue down these columns. At first glance, there appears to be no pattern in electron affinity across a row of the periodic table, as shown in the figure below. When these data are listed along with the electron configurations of these elements, however, they make sense.

These data can be explained by noting that electron affinities are much smaller than ionization energies. As a result, elements such as helium, beryllium, nitrogen, and neon, which have unusually stable electron configurations, have such small affinities for extra electrons that no energy is given off when a neutral atom of these elements picks up an electron. These configurations are so stable that it actually takes energy to force one of these elements to pick up an extra electron to form a negative ion.

There is no doubt that sodium reacts vigorously with chlorine to form NaCl. The only question is whether it is legitimate to assume that this reaction occurs because chlorine atoms "like" electrons more than sodium atoms.

The first ionization energy for sodium is one and one-half times larger than the electron affinity for chlorine. Thus, it takes more energy to remove an electron from a neutral sodium atom than is given off when the electron is picked up by a neutral chlorine atom. We will obviously have to find another explanation for why sodium reacts with chlorine to form NaCl.

Before we can do this, however, we need to know more about the chemistry of ionic compounds. Practice Problem 3: Use the Bohr model to calculate the wavelength and energy of the photon that would have to be absorbed to ionize a neutral hydrogen atom in the gas phase. Click here to check your answer to Practice Problem 3 Click here to see a solution to Practice Problem 3. Practice Problem 4: Predict which element in each of the following pairs has the larger first ionization energy.

Practice Problem 5: Predict the group in the periodic table in which an element with the following ionization energies would most likely be found. Practice Problem 6: Use the trends in the ionization energies of the elements to explain the following observations. Skip to main content.

Periodic Properties. Search for:. Ionization Energy. Learning Objective Recognize the general periodic trends in ionization energy. Key Points The ionization energy is the energy required to remove an electron from its orbital around an atom to a point where it is no longer associated with that atom.

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 therefore are more tightly bound harder to remove. Show Sources Boundless vets and curates high-quality, openly licensed content from around the Internet.

Licenses and Attributions. CC licensed content, Shared previously. So, with that out of the way, let's think about how hard it will be ionize different elements in the periodic table.

In particular, how hard it is to turn them into cations. And to think about that, we'll introduce an idea called ionization energy. Ionization energy And this is defined, this is defined as the energy required, energy required to remove an electron, to remove an electron. So, it could've even been called cationization energy because you really see energy required to remove an electron and make the overall atom more positive.

So, let's think about the trends. And we already have a little bit of background on the different groups of the periodic table. So, for example, if we were to focus on, especially we could look at group one, and we've already talked about how Hydrogen's a bit of a special case in group one but if we look at everything below Hydrogen. If we look at the Alkali, if we look at the Alkali metals here we've already talked about the fact that these are very willing to lose an electron.

Because if they lose an electron they get to the electron configuration of the noble gas before it. So, if Lithium loses an electron then it has an outer shell electron configuration of Helium. It has two outer electrons and that's kind of, we typically talk about the Octet Rule but if we're talking about characters like Lithium or Helium they're happy with two 'cause you can only put two electrons in that first shell.

But all the rest of 'em, Sodium, Potassium, etc. Lithium, if you remove an electron, it would get to Helium and it would have two electrons in its outer shell. So, you can imagine that the ionization energy right over here, the energy required to remove electrons from your Alkali Metals is very low. So, let me just write down this is So, when I say low, I'm talking about low ionization energy. Now, what happens as we move to the right of the periodic table?

In fact, let's go all the way to the right on the periodic table. Well, if we go here to the Noble Gases, the Noble Gases we've already talked about. They're very, very, very stable. They don't want no one, they don't want their electron configurations messed with.

So, it would be very hard Neon on down has their eight electrons that mumbling Octet Rule. Helium has two which is full for the first shell, and so it's very hard to remove an electron from here, and so it has a very high ionization energy.

Low energy, easy to remove electrons. Or especially the first electron, and then here you have a high ionization energy. I know you have trouble seeing that H. So, this is high, high ionization energy, and that's the general trend across the periodic table.

As you go from left to right, you go from low ionization energy to high ionization energy. Now, what about trends up and down the periodic table? Well, within any group, if we, even if we look at the Alkali, if we look at the Alkali Metals right over here, if we're down at the bottom, if we're looking at, if we're looking at, say, Cesium right over here, that electron in the, one, two, three, four, five, six, in the sixth shell, that's going to be further from that one electron that Lithium has and its second shell.

So, it's going to be, it's going to be further away.