For the SAT II Chemistry test, remember that electron affinity becomes more negative as we move across a period. This means that it’s easier to add an electron to elements, the farther to the right you travel on the periodic table. Why? Again, this is because the higher Zeff increases the nuclear attraction for the incoming electron. Important exceptions to this rule are the noble gases: He, Ne, Ar, Kr, and Xe. They have electron affinities that are positive (meaning very low), because if they were to accept another electron, that electron would have to go into a new, higher-energy subshell, and this is energetically unfavorable.
Electron affinities do not change very much as you go down a group. This is because the lower electron-nucleus attraction that’s seen as we go down a group is pretty evenly counterbalanced by a simultaneous lowering in electron-electron repulsion. Remember that there is no clear trend for electron affinity as you go down a group on the periodic table—this fact could come up in a synthesis of knowledge question!
Electronegativity
Electronegativity is a measure of the attraction an atom has for electrons when it is involved in a chemical bond. Elements that have high ionization energy and high electron affinity will also have high electronegativity since their nuclei strongly attract electrons. Electronegativity increases from left to right as we move across a period and decreases as we move down any group or family.
By now, these trends should make sense. You know that ionization energies tend to decrease with increasing atomic number in a group, although there isn’t a significant change in electron affinity, so it makes sense that atoms’ attraction for electrons in a bond would also increase as their Zeff increased. We will discuss the concept of electronegativity further in the next section, when we discuss chemical bonding.
Here’s a summary of the trends we discussed in this section. Make sure to memorize them!
