In this tutorial about electron affinity, we will cover its definition, relevant periodic table trends, and factors that influence it.
Topics Covered in Other Articles
- Ionization Energy
- Periodic Table Trends
- Lattice Energy
- Electron Orbitals & Orbital Shapes
What is Electron Affinity?
Chemists define electron affinity as the change in energy, measured in units of kJ/mole, experienced by a gaseous atom when an electron joins it. This electron typically appears in the form of a negative ion. Electron affinity differs from electronegativity, which we define as the ability of an atom to attract an electron toward itself.
We tend to liken electron affinity to an atom’s “likelihood,” or “chance,” of gaining an electron. It is the opposite of ionization energy, the energy required to ionize a gaseous atom and consequently remove an electron. Essentially, electron affinity pertains to the energy changes that accompany the gain of one electron, and ionization energy those that accompany losing one electron.
Although ionization energies always involve the formation of positive ions, electron affinity energies describe the generation of negative ions.
First and Second Electron Affinities
There are two types of electron affinity, first and second. First electron affinity involves the addition of an electron to a neutral atom. Because this exothermic process releases energy, first electron affinities are negative values.
Second electron affinity pertains to the addition of an electron to a negative ion. This endothermic process requires enough energy to eclipse the release of energy from the electron attachment process. This then results in positive second electron affinity values.
Electron Affinity Trends
As we travel to the right across periods, electron affinities become more exothermic. Scientists attribute this pattern to the addition of electrons closer to the nuclei of these more rightward atoms.
As elements trend to more to the right, the closer added electrons sit to their nuclei. These electrons exhibit a stronger attraction to the nuclei as a result of this proximity, explaining the exothermic nature of their electron affinities.
As we travel down groups, electron affinities become less exothermic. Electron proximity to these respective nuclei also influences this phenomenon. But contrary to those of the previous trend, these electrons are placed in higher energy levels. This means these more downward elements contain more electrons further from their nuclei.
Scientists explain this by referencing the lack of attraction between these nuclei and electrons further from them. Atoms of elements exhibiting this type of interaction release less energy upon the addition of an extra electron, explaining the decreased exothermic nature of their electron affinities.
Below is a visual representation of electron affinity trends throughout the periodic table. As discussed, electron affinities increase from left to right across periods; electron affinities decrease from top to bottom down groups.
You may be wondering how an increase in valence electrons affects these electron affinity trends, specifically that involve groups on the periodic table. Since the amount of valence electrons grows as we move down each group, these lower elements should, presumably, exhibit more stability and higher electron affinities.
But this assumption conflicts with the group trend discussed above. As noted, electron affinity decreases, rather than increases, down groups. We can address the clash between concepts using the shielding effect. This rule acknowledges that while nuclei attract valence electrons, repulsion forces counter this attraction. These repulsion forces are generated by inner electrons positioned between the nuclei and the outer electrons.
As we travel down each group, the influence of the shielding effect increases. This prompts a simultaneous decrease in the attraction between outer electrons and their respective nuclei, producing elements with less electron affinity.
Exception of Fluorine
Fluorine presents another caveat regarding the group electron affinity trend. Although first electron affinities generally decrease as we travel down a group, corresponding to less energy involvement in the formation of negative ions, fluorine atoms break this pattern.
Given its position at the top of its respective group, you would expect fluorine to exhibit a relatively high electron affinity. Due to fluorine’s very small atomic radius, the space surrounding its nucleus is also very small, thereby increasing the attraction between the incoming electron and the fluorine nucleus–and, by extension, its electron affinity.
However, we must keep in mind that the incoming electron would be entering a crowded area already affected by high levels of repulsion. As we talked about, repulsion and the shielding effect decreases the attraction between electrons and the nucleus, consequently lessening electron affinity. These repulsion effects are strong in fluorine atoms, prompting them to display smaller-than-expected electron affinities.
Metals vs Nonmetals
Generally, metals possess lower electron affinities while nonmetals have higher ones.
Electron Affinity of Metals
Metals want to form stable octets via the formation of cations; to accomplish this, they tend to give up valence electrons. They absorb energy when they lose electrons, contributing to lower, and endothermic, affinities.
Electron Affinity of Nonmetals
Conversely, nonmetals like to gain electrons to form anions as they pursue a full octet. They release energy when they take on electrons, producing higher, and exothermic, electron affinities.
What Affects the Electron Affinity of an Atom?
Once again, when we talk about electron affinity, we reference the change in energy an atom experiences when it gains an electron. This process reflects the amount of attraction between this incoming electron and the nucleus of the involved atom. The stronger the attraction, the more energy is released—and the higher the electron affinity for that atom will be.
The two main factors that influence these trends, as seen in our examples, include atomic size and nuclear charge. Nuclei possessing greater charges attract electrons more strongly, resulting in larger electron affinities. Conversely, less potent nuclear charges facilitate smaller electron affinities.
Regarding atomic size, smaller atoms offer less space for electrons to gather, including the incoming electron. As a result, this extra electron would position itself closer to the nucleus than it would in a larger atom. This results in larger electron affinity values for smaller atoms as a result of the increased attraction between this incoming electron and the nucleus. On the other hand, larger atoms tend to display smaller electron affinities because they offer more space for electrons to share with themselves as well as the incoming electron.
But we must not underestimate the effects of repulsion and shielding on electron affinity. For instance, smaller atoms may exhibit more attraction and larger electron affinities, but their lack of space for gathering electrons also produces increased repulsion between these particles. Repulsion lessons the attraction between the incoming electron and the nucleus, facilitating lowered electron affinities. The shielding effect acts similarly.
We must consider each of these factors carefully, as their effects will vary based on the characteristics of each element. Investigating the questions below will provide an example of this analysis.
Which Element Has the Lowest Electron Affinity?
Mercury has the lowest electron affinity among the elements. This status can be attributed to its delineation as a metal. Metals are, generally, more likely to lose electrons than gain them in their pursuit of a full, stable octet. Mercury has a relatively large atomic radius, which also contributes to its low value.
Which Element Has the Highest Electron Affinity?
Chlorine has the highest electron affinity among the elements. Its high electron affinity can be attributed to its large atomic radius, or size. Because chlorine’s outermost orbital is 3p, its electrons have a large amount of space to share with an incoming electron. This minimizes repulsions between these particles to a degree that outweighs the negative effects of its large size on attraction.