# High School Chemistry/Electron Affinity

The final periodic trend for our discussion is electron affinity. We have talked about atomic structure, electronic configurations, size of the atoms and ionization energy. And now, the final periodic trend we will study is how an atom can gain an electron and the trends that exist in the Periodic Table.

## Lesson Objectives

* Define electron affinity.
* Describe the trend for electron affinity on the Periodic Table.

## The Energy Process When an Electron is Added to an Atom

Atoms can gain or lose electrons. When an atom gains an electron, energy is given off and is known as the electron affinity. Electron affinity is defined as the energy released when an electron is added to a gaseous atom or ion.
    ${\text{T}}+e^{-}\rightarrow {\text{T}}^{-}\,\!$
When most reactions occur that involve the addition of an electron to a gaseous atom, potential energy is released.
    ${\text{Br}}+e^{-}\,\!$$\rightarrow \,\!$${\text{Br}}^{1-}\,\!$ $[{\text{Ar}}]4s^{2}4p^{5}\,\!$ $[{\text{Ar}}]4s^{2}4p^{6}\,\!$
Let's look at the electron configurations of a few of the elements and the trend that develops within groups and periods. Take a look at Table 10.10, the electron affinity for the Halogen family.
Table 10.10: Electron Affinities for Group 17 :
| Element | Electron Configuration | Electron Affinity, kJ/mol |
| Fluorine (F) | [He]2s$^2$2p$^5$ | −328 |
| Chlorine (Cl) | [Ne]3s$^2$3p$^5$ | −349 |
| Bromine (Br) | [Ar]4s$^2$4p$^5$ | −325 |
| Iodine (I) | [Kr]5s$^2$5p$^5$ | −295 |
As you can see, the electron affinity generally decreases (becomes less negative) going down a group because of the increase in size of the atoms. Remember that the atoms located within a family but lower on the periodic table are larger since there are more electrons filling more energy levels. For example, an atom of chlorine is smaller than iodine; or, an atom of oxygen is smaller than sulfur. When an electron is added to a large atom, less energy is released because the electron cannot move as close to the nucleus as it can in a smaller atom. Therefore, as the atoms in a family get larger, the electron affinity gets smaller.
There is an exception to this when it involves certain small atoms. Electron affinity for fluorine is less than chlorine most likely due to the electron-electron repulsions that occur between the electrons where n = 2. This phenomenon is observed in other families as well. For instance, the electron affinity for oxygen is less than the electron affinity for sulfur. Electron affinity of all of the elements in the second period is less than the ones below them due to the fact that the elements in the second period have such small electron clouds that electron repulsion is greater than that of the rest of the family.

## Nonmetals Tend to Have the Highest Electron Affinity

Overall, the Periodic Table shows the general trend similar to the one below.
Electron Affinities for Period 4 Main Group Elements :
| Element | Electron Configuration | Electron Affinity |
| Potassium (K) | [Ar]4s$^1$ | −48 kJ/mol |
| Calcium (Ca) | [Ar]4s$^2$ | −2.4 kJ/mol |
| Gallium (Ga) | [Ar]4s$^2$4p$^1$ | −29 kJ/mol |
| Germanium (Ge) | [Ar]4s$^2$4p$^2$ | −118 kJ/mol |
| Arsenic (As) | [Ar]4s$^2$4p$^3$ | −77 kJ/mol |
| Selenium (Se) | [Ar]4s$^2$4p$^4$ | −195 kJ/mol |
| Bromine (Br) | [Ar]4s$^2$4p$^5$ | −325 kJ/mol |
| Krypton (Kr) | [Ar]4s$^2$4p$^6$ | 0 kJ/mol |
The general trend in the electron affinity for atoms is almost the same as the trend for ionization energy. This is because both electron affinity and ionization energy are highly related to atomic size. Large atoms have low ionization energy and low electron affinity. Therefore, they tend to lose electrons and do not tend to gain electrons. Small atoms, in general, are the opposite. Since they are small, they have high ionization energies and high electron affinities. Therefore, the small atoms tend to gain electrons and tend not to lose electrons. The major exception to this rule are the noble gases. They are small atoms and do follow the general trend for ionization energies. The noble gases, however, do not follow the general trend for electron affinities. Even though the noble gases are small atoms, their outer energy levels are completely filled with electrons and therefore, an added electron cannot enter their outer most energy level. Any electrons added to a noble gas would have to be the first electron in a new (larger) energy level. This causes the noble gases to have essentially zero electron affinity. This concept is discussed more thoroughly in the next chapter.
When atoms become ions, the process involves either the energy released through electron affinity or energy being absorbed with ionization energy. Therefore, the atoms that require a large amount of energy to release an electron will most likely be the atoms that give off the most energy while accepting an electron. In other words, non-metals will most easily gain electrons since they have large electron affinities and large ionization energies; and, metals will lose electrons since they have the low ionization energies and low electron affinities.
Now let's add this last periodic trend to our Periodic Table representation and our periodic trends are complete.
The development and arrangement of the Periodic Table of Elements is examined (free registration required). Video on Demand – The World of Chemistry – The Periodic Table.

## Lesson Summary

* Electron affinity is the energy required (or released) when an electron is added to a gaseous atom or ion. Electron affinity generally increases going up a group and increases left to right across a period.
* Non-metals tend to have the highest electron affinities.

## Review Questions

* Define electron affinity and show an example equation.
* Choose the element in each pair that has the lower electron affinity:
    (a) Li or N
    (b) Na or Cl
    (c) Ca or K
    (d) Mg or F
* Why is the electron affinity for calcium much higher than that of potassium?
* Draw a visual representation of the periodic table describing the trend of electron affinity.
* Which of the following would have the largest electron affinity?
    (a) Se
    (b) F
    (c) Ne
    (d) Br
* Which of the following would have the smallest electron affinity?
    (a) Na
    (b) Ne
    (c) Al
    (d) Rb
* Place the following elements in order of increasing electron affinity: Te, Br, S, K, Ar.
* Place the following elements in order of decreasing electron affinity: S, Sn, Pb, F, Cs.
* Describe the trend that would occur for electron affinities for elements in Period 3. Are there any anomalies? Explain.
* Comparing the electron affinity (EA) of sulfur, S, and phosphorus, P:
    (a) S has a higher EA because its radius is smaller.
    (b) P has a higher EA because its radius is smaller.
    (c) S has a higher EA because its p sub-shell is half full.
    (d) P has a higher EA because its p sub-shell is half full.
    (e) they have the same EA because they are next to each other in the Periodic Table.

## Vocabulary
### electron affinity

    The energy required to add an electron to a gaseous atom or ion.
← Ionization Energy · Ions and the Compounds They Form →
This material was adapted from the original CK-12 book that can be found here. This work is licensed under the Creative Commons Attribution-Share Alike 3.0 United States License