The textbook definition says that electron affinity is \"the energy change assoc

Question

The textbook definition says that electron affinity is "the energy change associated with the addition of an electron to a gaseous atom."

My textbook also says, "When we go down a group, electron affinity should become more positive (less energy released), since the electron is added at increasing distances from the nucleus."

I understand that the atoms become larger as you go down a group. What I don't understand is what it means to have less energy released. Does this mean that it's easier to add an electron to a larger atom because it requires less energy for a negative ion to form?

I thought that it would be more difficult to add an electron to a larger atom because the core atoms shield the nucleus' positive charge. Thus since the nuclear charge is weaker in a larger atom, any electron coming near it won't want to bond as readily.

Thanks for your clarifications.

Explanation / Answer

Electron affinity decreases down the groups and from right to left across the periods on the periodic table because the electrons are placed in a higher energy level far from the nucleus, thus a decrease from its pull. However, one might think that since the number of valence electrons increase going down the group, the element should be more stable and have higher electron affinity. One fails to account for the shielding affect. As one goes down the period, the shielding effect increases, thus repulsion occurs between the electrons. This is why the attraction between the electron and the nucleus decreases as one goes down the group in the periodic table.

The electron affinity is a measure of the attraction between the incoming electron and the nucleus - the stronger the attraction, the more energy is released

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