Order the atoms in each of the following sets from the least exothermic electron affinity to the most. a. \(\mathrm{S}, \mathrm{Se}\) b. \(\mathrm{F}, \mathrm{Cl}, \mathrm{Br}, \mathrm{I}\)

Understanding how various properties change across the periodic table allows us to predict and explain the chemical behavior of elements. One such property is electron affinity, which typically becomes more exothermic as we move from left to right across a period. This is due to the increasing nuclear charge, which enhances the attraction of electrons. However, there are exceptions, such as the trend not being uniformly gradual among different groups.

Similarly, there are vertical trends to consider. For instance, within groups, electron affinity tends to become less exothermic as we move down the group, owing to the increasing distance from the nucleus and the rising effect of electron shielding, making it easier for the atom to accommodate additional electrons.

Chemical properties of an element are largely dictated by the behavior of their electrons, particularly the valence electrons which are involved in bonding and reactions. Electron affinity, a key chemical property, measures how much energy is released when an electron is added to a neutral atom to form a negatively charged ion. A high exothermic electron affinity means that an atom releases more energy upon gaining an electron, which usually correlates with higher reactivity and stronger ability to attract electrons in chemical bonds.

The elements in Group 16 of the periodic table are often referred to as the 'chalcogens.' They include oxygen, sulfur, selenium, tellurium, and polonium. The exercise presents us with sulfur (S) and selenium (Se), which are directly above and below each other in this group. As we progress down the group, we see a decrease in electron affinity since additional electron shells mean the valence electrons are further from the nucleus and more shielded by inner electrons. This explains why sulfur has a more exothermic electron affinity than selenium.

Illustrative Comparison:

• Sulfur (S) - higher electron affinity, more exothermic.
• Selenium (Se) - lower electron affinity, less exothermic.

Group 17 is known as the halogens, which are highly reactive nonmetals including fluorine (F), chlorine (Cl), bromine (Br), and iodine (I). They have particularly high electron affinities due to their one-electron short of a noble gas electron configuration. This high electron affinity decreases as you go down the group - meaning fluorine has the most exothermic electron affinity, followed by chlorine, bromine, and iodine, respectively. Their reactivity is a notable chemical property that makes the halogens important in many chemical reactions.

Exothermic Trend in Halogens:

• Fluorine (F) - highest electron affinity, most exothermic.
• Chlorine (Cl) - lower than F.
• Bromine (Br) - lower than Cl.
• Iodine (I) - lowest of the group, least exothermic.

An exothermic process is one that releases energy to the surrounding environment, usually in the form of heat. Measuring the exothermicity of electron affinity is an evaluation of how much energy is released when an atom gains an electron. A greater release of energy results in a more exothermic electron affinity. This information is particularly useful in predicting the stability of ions formed from atoms, as those with more exothermic electron affinities tend to form more stable ions.

Key Points:

• An exothermic electron affinity reflects a spontaneous process.
• Atoms with higher exothermic electron affinities often form stronger ionic bonds.
• This property can influence the reactivity and compound formation behavior of elements.