Question

Write the electron configuration for each ion. What do all of the electron configurations have in common? (a) \(\mathrm{F}^{-}\) (b) \(\mathrm{P}^{3-}\) (c) \(\mathrm{Li}^{+}\) (d) \(\mathrm{Al}^{3+}\)

Short answer

The electron configurations are \text{F}^{-}:\text{ }1s^2 2s^2 2p^6, \text{ }P^{3-}:\text{ }1s^2 2s^2 2p^6 3s^2 3p^6, \text{ }Li^{+}:\text{ }1s^2, \text{ }and \text{ }Al^{3+}:\text{ }1s^2 2s^2 2p^6. All configurations represent ions with a complete octet, similar to noble gases.

Step-by-step solution

Write Electron Configuration for \text{F}^{-}

Start with the neutral fluorine atom, which has an atomic number of 9. The electron configuration for neutral fluorine is \(1s^2 2s^2 2p^5\). Since \text{F}^{-} has gained one additional electron, it will be \(1s^2 2s^2 2p^6\).

Write Electron Configuration for \text{P}^{3-}

Start with the neutral phosphorus atom, which has an atomic number of 15. The electron configuration is \(1s^2 2s^2 2p^6 3s^2 3p^3\). For \text{P}^{3-}, three extra electrons are added making the configuration \(1s^2 2s^2 2p^6 3s^2 3p^6\).

Write Electron Configuration for \text{Li}^{+}

Neutral lithium has an atomic number of 3, with the configuration \(1s^2 2s^1\). In the \text{Li}^{+} ion, one electron is lost, resulting in the configuration \(1s^2\).

Write Electron Configuration for \text{Al}^{3+}

Aluminum has an atomic number of 13, and the neutral atom configuration is \(1s^2 2s^2 2p^6 3s^2 3p^1\). For \text{Al}^{3+}, three electrons are lost, two from the 3p orbital and one from the 3s, leaving us with \(1s^2 2s^2 2p^6\).

Identify Common Feature of the Configurations

Each ion has a complete octet, meaning that the \(s\) and \(p\) orbitals in their outermost shell are fully occupied with no unpaired electrons. This configuration is similar to that of a noble gas.

Key concepts

Electron Configuration

Electron configuration is a fundamental concept in chemistry that describes the arrangement of electrons in an atom's orbitals. Atoms seek the most stable structure possible, and electrons fill orbitals in a way that follows specific rules known as the Aufbau principle, Pauli exclusion principle, and Hund's rule.

For example, fluorine has an atomic number of 9, meaning it has 9 protons and, when neutral, 9 electrons. Its electron configuration is written as \(1s^2 2s^2 2p^5\). The superscripts represent the number of electrons in each orbital, and the numbers before the s and p refer to the energy level, or shell, where these orbitals are located. The same principles apply when determining the electron configuration for ions, adding or removing electrons depending on the charge.

When we take a fluorine atom and add an extra electron to form the \(\mathrm{F}^{-}\) ion, the additional electron fills the remaining space in the 2p orbital, resulting in \(1s^2 2s^2 2p^6\), which is a more stable, lower-energy configuration.

Valence Electrons

Valence electrons play a pivotal role in chemical reactions, as these are the electrons found in the outermost shell of an atom and thus the ones involved in forming chemical bonds. The number of valence electrons determines the reactivity of an element and dictates how an atom will bond with others.

In neutral atoms, the number of valence electrons is the same as the group number for the elements in the main group of the periodic table. For instance, aluminum is in group 13 and has three valence electrons. However, when it forms ions like \(\mathrm{Al}^{3+}\), it loses three electrons and ends up with a configuration resembling the nearest noble gas, neon, with eight valence electrons in the previous shell.

Octet Rule

The octet rule is a chemical rule of thumb that reflects observation that atoms of main-group elements tend to bond in such a way that each atom has eight electrons in its valence shell, giving it the same electron configuration as a noble gas. The rule is applicable to the main-group elements, especially carbon, nitrogen, oxygen, and the halogens, but there are exceptions.

This rule explains why the ions in the exercise - \(\mathrm{F}^{-}\), \(\mathrm{P}^{3-}\), \(\mathrm{Li}^{+}\), and \(\mathrm{Al}^{3+}\) - have adopted electron configurations that satisfy the octet rule. By gaining or losing electrons, these ions achieve a full valence shell resembling that of a noble gas, which is energetically favorable.

Ionic Charge

Ionic charge is determined by the loss or gain of valence electrons, which ultimately affects the electron configuration of the element. Ions form when atoms gain or lose electrons to achieve a full octet, resulting in either a positive charge (cation) when electrons are lost, or a negative charge (anion) when electrons are gained.

In our exercise, \(\mathrm{Li}^{+}\) is a cation which has lost an electron, while \(\mathrm{F}^{-}\) and \(\mathrm{P}^{3-}\) are anions having gained one and three electrons respectively. Aluminum becomes a cation \(\mathrm{Al}^{3+}\) by losing three electrons. The ionic charge influences not only the electron configuration but also the chemical and physical properties of the ion.