Question

Write general outer electron configurations \(\left(n s^{x} n p^{y}\right)\) for groups \(6 \mathrm{~A}\) and \(7 \mathrm{~A}\) in the periodic table. The electron affinity of each group 7A element is more negative than that of each corresponding group 6A element. Use the electron configurations to explain why this is so.

Short answer

Group 6A elements have the outer electron configuration of \((n-1)d^{10}ns^2np^4\), while Group 7A elements have \((n-1)d^{10}ns^2np^5\). Group 7A elements have more negative electron affinities because they are closer to having a complete octet, making them more energetically favorable to gain an additional electron.

Step-by-step solution

Writing electron configurations for Group 6A

The elements in Group 6A (the oxygen family) have electron configurations that end in \(ns^2np^4\). These elements have six electrons in their outermost shell: two in the \(s\) orbital and four in the \(p\) orbitals. This configuration can be represented as \((n-1)d^{10}ns^2np^4\), where \(n\) corresponds to the period of the element.

Writing electron configurations for Group 7A

The elements in Group 7A (the halogens) have electron configurations that end in \(ns^2np^5\). These elements have seven electrons in their outermost shell: two in the \(s\) orbital and five in the \(p\) orbitals. This configuration can be represented as \((n-1)d^{10}ns^2np^5\), where \(n\) again corresponds to the period of the element.

Explaining higher electron affinity in Group 7A

Electron affinity increases as the number of electrons in the outer shell approaches a full octet because the elements are more eager to accept an additional electron to achieve a stable noble gas configuration. Group 7A elements are one electron away from a full \(p\) orbital (a complete octet) and are thus more likely to accept an additional electron with greater attraction (more negative electron affinity) compared to Group 6A elements, which require two additional electrons to achieve a full octet.

Key concepts

Electron Affinity

Electron affinity refers to the amount of energy released when an electron is added to a neutral atom in the gaseous phase to form a negative ion. In essence, it is a measure of an atom's desire to gain an electron. Generally, high electron affinities indicate atoms that readily accept electrons, contributing to their higher reactivity.

As an example, elements in Group 7A of the periodic table, known as the halogens, exhibit significantly high electron affinities. This high electron affinity is because they are just one electron short of reaching the very stable configuration of a noble gas, which has a complete outer electron shell. In comparison, atoms in Group 6A require two electrons to achieve a full octet, leading to somewhat lower electron affinities.

Periodic Table Groups

The periodic table is organized into 18 vertical columns known as groups or families. Each group consists of elements that have similar physical and chemical properties due partly to their consistent outer electron configurations. As we move from left to right across a period (a horizontal row on the periodic table), the number of electrons in the outer shell of the elements increases.

For instance, Group 6A, also known as the chalcogens or oxygen family, includes elements such as oxygen, sulfur, and selenium, while Group 7A contains the halogens like fluorine and chlorine. Each group's elements tend to exhibit trends in reactivity, electron affinity, and other properties that are useful for predicting behavior in chemical reactions.

Outer Electron Shells

The outer electron shell of an atom is the highest energy level occupied by electrons and plays a crucial role in chemical bonding and reactivity. Electrons within these outer shells are called valence electrons and are responsible for an element's ability to form bonds with other atoms. An atom's chemical properties are largely determined by the number and configuration of its valence electrons.

For example, Group 6A elements have electron configurations ending in \(ns^2np^4\), which implies that they have six valence electrons—two in the \(s\) orbital and four in the \(p\) orbitals. On the other hand, Group 7A elements' electron configurations end in \(ns^2np^5\), indicating seven valence electrons. These configurations influence the electron affinity and reactivity of the elements.

Halogens

Halogens make up Group 7A of the periodic table and comprise elements such as fluorine, chlorine, bromine, iodine, and astatine. These nonmetals are characterized by their highly reactive nature and a strong tendency to form salts with metals. Halogens have seven valence electrons; their outer electron configuration is typically expressed as \(ns^2np^5\).

Because halogens are one electron short of a complete octet, they exhibit high electron affinities and react vigorously, especially with alkali and alkaline earth metals, to achieve a stable electron configuration. The formation of ionic compounds like sodium chloride (table salt) is a direct result of the halogens' desire to obtain that extra electron and satisfy the octet rule.

Oxygen Family

The oxygen family, or Group 6A, includes oxygen, sulfur, tellurium, and polonium, amongst others. These elements are known for their diverse range of oxidative states and compounds they form. The outer electron configuration of this family is \(ns^2np^4\), which means these elements have six valence electrons.

Members of the oxygen family are two electrons short of completing their octet, making them fairly reactive. While oxygen is essential for respiration and is the most abundant element in the Earth's crust, other members such as sulfur are utilized in the manufacture of fertilizers and various chemical processes. Their reactivity, while less than that of the halogens, is crucial to many biochemical and industrial reactions.