Question

Generalize The outer-electron configurations of elements in group 1 can be written as \(n \mathrm{s}^{1}\) , where \(n\) refers to the element's period and its principal energy level. Develop a similar notation for all the other groups of the representative elements.

Short answer

The outer-electron configurations for all the representative elements can be written as follows: - Group 1 (Alkali Metals): \(n\mathrm{s}^{1}\) - Group 2 (Alkaline Earth Metals): \(n\mathrm{s}^{2}\) - Group 13 (Boron Group): \(n\mathrm{s}^{2}\mathrm{p}^{1}\) - Group 14 (Carbon Group): \(n\mathrm{s}^{2}\mathrm{p}^{2}\) - Group 15 (Nitrogen Group): \(n\mathrm{s}^{2}\mathrm{p}^{3}\) - Group 16 (Chalcogens): \(n\mathrm{s}^{2}\mathrm{p}^{4}\) - Group 17 (Halogens): \(n\mathrm{s}^{2}\mathrm{p}^{5}\) - Group 18 (Noble Gases): \(n\mathrm{s}^{2}\mathrm{p}^{6}\) (except for Helium: \(1\mathrm{s}^{2}\))

Step-by-step solution

Group 2: Alkaline Earth Metals

The outer-electron configuration for group 2 elements can be written as: \(n\mathrm{s}^{2}\)

Group 13: Boron Group

The outer-electron configuration for group 13 elements can be written as: \(n\mathrm{s}^{2}\mathrm{p}^{1}\)

Group 14: Carbon Group

The outer-electron configuration for group 14 elements can be written as: \(n\mathrm{s}^{2}\mathrm{p}^{2}\)

Group 15: Nitrogen Group

The outer-electron configuration for group 15 elements can be written as: \(n\mathrm{s}^{2}\mathrm{p}^{3}\)

Group 16: Chalcogens

The outer-electron configuration for group 16 elements can be written as: \(n\mathrm{s}^{2}\mathrm{p}^{4}\)

Group 17: Halogens

The outer-electron configuration for group 17 elements can be written as: \(n\mathrm{s}^{2}\mathrm{p}^{5}\)

Group 18: Noble Gases

The outer-electron configuration for group 18 elements can be written as: \(n\mathrm{s}^{2}\mathrm{p}^{6}\) (except for Helium which only has the configuration \(1\mathrm{s}^{2}\)) Now we have generalized the outer-electron configurations for all the representative elements (groups 1, 2, and 13-18).

Key concepts

Periodic Table Groups

The periodic table is organized into vertical columns known as groups. Each group contains elements that exhibit similar chemical and physical properties. This similarity arises because the elements in a group share the same number of valence electrons, which are the electrons available for bonding. The main groups of the periodic table are often referred to as representative elements and include the s and p blocks. These are groups 1 and 2 (the s-block), and groups 13 through 18 (the p-block), displaying a regular pattern in their electron configurations as you move across periods.

Understanding periodic table groups is crucial because it allows scientists to predict the behavior of unknown elements based on their position. This organization also helps chemists understand the trends in properties like atomic size, ionization energy, and electronegativity across different groups.

Alkaline Earth Metals

Alkaline Earth Metals belong to Group 2 of the periodic table. These metals include beryllium, magnesium, calcium, strontium, barium, and radium. The defining feature of all the elements in this group is their outer electron configuration of \(n\mathrm{s}^{2}\), which signifies two electrons in the outermost shell.

These metals are known for being shiny and somewhat reactive, especially with water, though less so than the Group 1 Alkali Metals. They can lose their two outermost electrons to form ions with a +2 charge, making them important in forming various compounds.

• They play crucial roles in nature and industry, with calcium being essential for biological processes and magnesium widely used in materials.
• Their tendency to lose two electrons makes them useful in a variety of chemical reactions.

Alkaline Earth Metals demonstrate increasing reactivity and lower melting points moving down the group, making this knowledge useful for predicting their behavior.

Boron Group

The Boron Group comprises Group 13 of the periodic table. This group includes boron, aluminum, gallium, indium, and thallium. The elements in Group 13 possess an outer electron configuration of \(n\mathrm{s}^{2}\mathrm{p}^{1}\).

Boron, the first element in the group, is unique as it is a non-metal, whereas the rest are metals. Their electronic configuration leads these elements to frequently form covalent bonds. The Boron Group exhibits interesting behaviors:

• Boron is known for its hard, yet brittle nature and high melting point, and finds use in borosilicate glass and detergents.
• Aluminum is lightweight and highly malleable, used extensively in packaging and structural applications due to its reactivity being reduced by forming a protective oxide layer.

The Boron Group's distinct behavior showcases the breadth of properties across the periodic table, influenced by electron configurations.

Noble Gases

Noble Gases are located in Group 18 of the periodic table. This group includes helium, neon, argon, krypton, xenon, and radon. They are characterized by their complete outer electron shells, generally represented by the electron configuration \(n\mathrm{s}^{2}\mathrm{p}^{6}\). However, helium is an exception, having a configuration of \(1\mathrm{s}^{2}\).

The complete electron shells render Noble Gases remarkably stable and inert, exhibiting very low chemical reactivity. This characteristic is why they are sometimes referred to as "inert gases."

• They are used in lighting, welding, and as preservatives in food packaging due to their non-reactivity.
• Helium, being lighter than air, is famously used in balloons and has critical applications in cryogenics.

Noble Gases' properties make them invaluable in applications requiring non-reactive environments, benefiting a wide array of scientific and industrial uses due to their stable electron configurations.