Question

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

Short answer

For Set A (S, Se), the order from least negative to most negative electron affinity is: S, Se. For Set B (F, Cl, Br, I), the order is: I, Br, Cl, F.

Step-by-step solution

Find the position of S and Se in the periodic table

Sulfur (S) and Selenium (Se) are both in Group 16 of the periodic table. Sulfur is in Period 3, while Selenium is in Period 4.

Order the elements by electron affinity

Since Sulfur (S) is above Selenium (Se) in the same group, S will have a less negative electron affinity than Se. Thus, the order is: S, Se. For Set B: F, Cl, Br, I

Find the position of F, Cl, Br, and I in the periodic table

Fluorine (F), Chlorine (Cl), Bromine (Br), and Iodine (I) are all in Group 17 of the periodic table - the halogens. Fluorine is in Period 2, Chlorine is in Period 3, Bromine is in Period 4, and Iodine is in Period 5.

Order the elements by electron affinity

Since electron affinity decreases down a group, we can order the elements from least negative to most negative electron affinity as follows: I, Br, Cl, F.

Key concepts

Understanding the Periodic Table

The periodic table is a powerful tool for chemists, summarizing a wealth of information about the elements in a compact format. It is organized into rows (periods) and columns (groups), with elements arranged in order of increasing atomic number. Elements in the same group share similar chemical properties because they have the same number of electrons in their outermost shell.

For instance, when considering electron affinities - the energy change that occurs when an electron is added to a neutral atom - the periodic table reveals patterns. Generally, as we move from left to right across a period, electron affinity becomes more negative because the added electron is entering an orbital closer to the nucleus. Implementing this knowledge helps us understand why within a group, such as Group 16, sulfur (S) has a less negative electron affinity than selenium (Se), since S is located in the third period, while Se is in the fourth.

When students study how to predict and compare the electron affinities of different elements, recognizing the layout and organizing principles of the periodic table is essential. By understanding this, students grasp why elements in the same period increase in electron affinity from left to right and why groups show a decrease in electron affinity as you descend the group.

Halogens and Their Properties

Halogens are a group of highly reactive nonmetals located in Group 17 of the periodic table. This family consists of fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At). One of their most notable features is their electron affinity. Electron affinity measures an atom's tendency to accept an electron, releasing energy in the process.

The halogens are known for having a high electron affinity because they are one electron short of a full valence shell, which makes gaining an additional electron highly favorable. Fluorine has the highest electron affinity in the group, as it is the most electronegative element. As you move down the group, from fluorine to astatine, the electron affinity generally decreases. This is because the added electron will be farther from the nucleus and more shielded by inner electrons, thus less attracted to the atomic nucleus.

When students compare the electron affinities of halogens, they should remember this trend of decreasing electron affinity down the group. This explains the ordering of I, Br, Cl, and F, with iodine having the least negative electron affinity and fluorine having the most.

Chemical Periodicity and Trends

Chemical periodicity refers to the recurring trends and patterns in properties of elements that are visible within the periodic table. These patterns arise due to the periodic nature of the electrons' arrangement within atoms, particularly in the outer shells. For instance, the electron affinity - alongside other properties like atomic size, ionization energy, and electronegativity - changes in a predictable way across the periods and down the groups of the periodic table.

Within a period, electron affinity generally becomes more negative as the group number increases. This is because the nuclear charge increases with each added proton, drawing the added electron in more strongly. However, there are exceptions to this smooth trend due to the subshell configurations.

Down a group, the outermost electrons are farther from the nucleus with each successive element, resulting in a less negative electron affinity. This accounts for the trend observed with the halogens, as well as the elements of Group 16, where sulfur and selenium display decreasing electron affinity with increasing atomic number. Understanding chemical periodicity allows students to make educated predictions about an element's behavior and to effectively group elements with similar properties.