Question

Write the electron configuration for each of the following ions, and determine which ones possess noble-gas configurations: (a) \(\mathrm{Sr}^{2+}\), (b) \(\mathrm{Ti}^{2+}\), (c) \(\mathrm{Se}^{2-}\),(d) \(\mathrm{Ni}^{2+}\),(e) \(\mathrm{Br}^{-}\), (f) \(\mathrm{Mn}^{3+}\).

Short answer

The electron configurations for the given ions are: (a) \(\mathrm{Sr}^{2+}\): [Kr] - Noble gas configuration (b) \(\mathrm{Ti}^{2+}\): [Ar]3d^2 - Not a noble gas configuration (c) \(\mathrm{Se}^{2-}\): [Kr] - Noble gas configuration (d) \(\mathrm{Ni}^{2+}\): [Ar]3d^6 - Not a noble gas configuration (e) \(\mathrm{Br}^{-}\): [Kr] - Noble gas configuration (f) \(\mathrm{Mn}^{3+}\): [Ar]3d^2 - Not a noble gas configuration \(\mathrm{Sr}^{2+}\), \(\mathrm{Se}^{2-}\), and \(\mathrm{Br}^{-}\) possess noble-gas configurations.

Step-by-step solution

1. Write the electron configuration of the neutral atom

First, we should write the electron configuration of the neutral atom for each element. You can use the periodic table to find the order in which subshells are filled. (a) Sr: [Kr]5s^2 (b) Ti: [Ar]4s^2 3d^2 (c) Se: [Ar]4s^2 3d^{10} 4p^4 (d) Ni: [Ar]4s^2 3d^8 (e) Br: [Ar]4s^2 3d^{10} 4p^5 (f) Mn: [Ar]4s^2 3d^5

2. Write the electron configuration of each ion

Now, taking into account the charge of each ion, we will adjust the number of electrons to match the charge. Subshells should be emptied according to the (n+l) rule, respecting Hund's rule. (a) \(\mathrm{Sr}^{2+}\): [Kr] (b) \(\mathrm{Ti}^{2+}\): [Ar]3d^2 (c) \(\mathrm{Se}^{2-}\): [Ar]4s^2 3d^{10} 4p^6 (d) \(\mathrm{Ni}^{2+}\): [Ar]3d^8-2 = [Ar]3d^6 (e) \(\mathrm{Br}^{-}\): [Ar]4s^2 3d^{10} 4p^5+1 = [Ar]4s^2 3d^{10} 4p^6 (f) \(\mathrm{Mn}^{3+}\): [Ar]4s^2 3d^5-3 = [Ar]3d^2

3. Compare the electron configurations with noble gases

Finally, we will compare the electron configurations of each ion to the electron configurations of noble gases. Noble gases have completely filled s and p orbitals in their highest energy level. (a) \(\mathrm{Sr}^{2+}\): [Kr] - Noble gas configuration (b) \(\mathrm{Ti}^{2+}\): [Ar]3d^2 - Not a noble gas configuration (c) \(\mathrm{Se}^{2-}\): [Ar]4s^2 3d^{10} 4p^6 - Noble gas configuration ([Kr]) (d) \(\mathrm{Ni}^{2+}\): [Ar]3d^6 - Not a noble gas configuration (e) \(\mathrm{Br}^{-}\): [Ar]4s^2 3d^{10} 4p^6 - Noble gas configuration ([Kr]) (f) \(\mathrm{Mn}^{3+}\): [Ar]3d^2 - Not a noble gas configuration So, \(\mathrm{Sr}^{2+}\), \(\mathrm{Se}^{2-}\), and \(\mathrm{Br}^{-}\) possess noble-gas configurations.

Key concepts

Noble Gas Configuration

A noble gas configuration is a key concept in understanding how ions stabilize themselves by achieving a completely filled electron shell, similar to the nearest noble gas. Noble gases are known for their stability, since their outermost electron shell is fully occupied, providing them with minimal reactivity. To achieve a noble gas configuration, atoms will gain, lose, or share electrons to fill or empty their outer electron shell. This process results in a stable electronic arrangement. For example:

• Strontium Ion (\(\mathrm{Sr}^{2+}\)): When it loses 2 electrons, it shares the same electron configuration as Krypton, \([\mathrm{Kr}]\).
• Selenium Ion (\(\mathrm{Se}^{2-}\)): Gains 2 electrons and matches Krypton, \([\mathrm{Kr}]\).
• Bromine Ion (\(\mathrm{Br}^{-}\)): By gaining 1 electron, it achieves the configuration of Krypton, \([\mathrm{Kr}]\).

Achieving a noble gas configuration is an important principle as it explains why elements form ions with certain charges. By doing so, they reach a low energy state, making them more stable compared to their neutral form.

Electron Subshells

Electron subshells describe regions where electrons are likely to be found in an atom. These regions are part of a larger structure known as electron shells, which are divided into subshells labeled as \(s\),\(p\),\(d\), and \(f\).In each subshell, electrons have specific levels of energy.Understanding subshells provides insight into why electrons occupy certain positions within an atom's structure. For instance:

• The \(s\) subshell can hold up to 2 electrons.
• The \(p\) subshell can accommodate up to 6 electrons.
• The\(d\) subshell has room for up to 10 electrons.
• The\(f\) subshell can contain up to 14 electrons.

By analyzing the electron configuration of ions, we determine electrons are removed from the highest energy subshell first. This often means electrons in the \(s\) subshell get removed before the\(d\) subshell in transition metals, as seen in the example of \(\mathrm{Ni}^{2+}\), where two outer electrons are lost from the \(4s\) subshell first.

Periodic Table

The periodic table is an essential tool in understanding electron configurations and the behavior of elements. It organizes elements in a way that highlights periodicity, meaning trends and patterns repeat throughout the table.Each element's position on the periodic table corresponds to its atomic structure, including:

• Electron configuration and energy levels, which determine chemical behavior and properties.
• The atomic number, indicating the number of protons and consequently the number of electrons in a neutral atom.

For effective electron configuration, elements are grouped into blocks:

• s-block: Includes Groups 1 and 2, where electron configurations end in \(s\) subshells.
• p-block: Holds Groups 13 to 18, with electron configurations ending in \(p\) subshells.
• d-block: Known as transition metals, associated with \(d\) subshells.
• f-block: Contains lanthanides and actinides, with electrons filling \(f\) subshells.

Utilizing the periodic table, we predict electron configurations and predict how ions will form. For instance, Selenium is located in the \(p\) block, indicating its outer electrons are in \(p\) orbitals. Understanding these patterns helps predict chemical reactions and behavior.