Question

Order the atoms in each of the following sets from the least exothermic electron affinity to the most. a. \(\mathrm{N}, \mathrm{O}, \mathrm{F}\) b. Al, Si, \(\mathrm{P}\)

Short answer

For set A, the order from the least exothermic electron affinity to the most is: \(N < O < F\). For set B, the order is: \(Al < Si < P\).

Step-by-step solution

Locating the atoms on the periodic table

Locate the elements nitrogen (N), oxygen (O), and fluorine (F) on the periodic table. They are all in the second period (row) and are in the following order: N, O, F.

Applying the electron affinity trend

Since all three elements are in the same period of the periodic table, we can apply the general trend of electron affinity increasing from left to right in the same period. This means that nitrogen (N) has the least exothermic electron affinity, followed by oxygen (O), and finally fluorine (F) with the most exothermic electron affinity.

Ordering the atoms based on electron affinity

Using the information from Steps 1 and 2, we can order the atoms in Set A from the least exothermic electron affinity to the most: N < O < F. #Set B: Al, Si, P#

Locating the atoms on the periodic table

Locate the elements aluminum (Al), silicon (Si), and phosphorous (P) on the periodic table. They are all in the third period (row) and are in the following order: Al, Si, P.

Applying the electron affinity trend

Since all three elements are in the same period of the periodic table, we can apply the general trend of electron affinity increasing from left to right in the same period. This means that aluminum (Al) has the least exothermic electron affinity, followed by silicon (Si), and finally phosphorous (P) with the most exothermic electron affinity.

Ordering the atoms based on electron affinity

Using the information from Steps 1 and 2, we can order the atoms in Set B from the least exothermic electron affinity to the most: Al < Si < P.

Key concepts

Periodic Table Trends

The periodic table is a valuable tool that allows us to quickly understand various chemical properties of elements. One such property is electron affinity, which is the energy change when an electron is added to a neutral atom. Generally, across a period (row), electron affinity becomes more exothermic from left to right. This is because atoms tend to desire more electrons to fill their outer electron shell as they approach the noble gases.
You can observe this trend within individual sets of elements. For instance, in the second period, nitrogen (N), oxygen (O), and fluorine (F) show increasing exothermicity in electron affinity. Likewise, in the third period, elements like aluminum (Al), silicon (Si), and phosphorus (P) also follow this increasing trend to the right. Nevertheless, certain exceptions and variations might occur, often due to electron-electron repulsions or the particular electronic configurations of the elements.
Recognizing these periodic table trends is crucial as they provide insight into how reactive an element might be or how it will interact chemically with other elements.

Exothermic Reactions

Exothermic reactions release energy to the surroundings, often in the form of heat. When you think about electron affinity as an exothermic process, it's important to understand that it involves an atom's tendency to gain an extra electron and the subsequent energy release.
In cases like fluorine, which has a high electron affinity, the process of gaining an electron releases a lot of energy due to the strong attraction between the added electron and the nucleus. This release of energy signifies the reaction is exothermic.

• Energy release is more significant in elements with high electron affinity.
• Conversely, certain elements may even have near-zero or slightly endothermic electron affinities.

Understanding electron affinities in terms of exothermic reactions helps us predict elemental reactivity and how these elements might engage in chemical processes.

Group and Period Trends

The periodic table not only allows us to identify trends across periods but also down groups. In general, across a period, elements become more nonmetallic and tend to gain electrons more readily, hence more exothermic electron affinities.
Within a group, as you move down, the increased number of electron shells causes added electrons to be further from the nucleus and shielded by inner electron layers. Thus, elements tend to have less exothermic or even endothermic electron affinities.

• Top to bottom in a group: less exothermic.
• Left to right in a period: more exothermic.

Therefore, these group and period trends give a comprehensive view of how and why electron affinity values change across the periodic table, impacting the chemical behavior and bonding tendencies of elements.