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 ion configurations are F-: 1s2 2s2 2p6, P3-: 1s2 2s2 2p6 3s2 3p6, Li+: 1s2, and Al3+: 1s2 2s2 2p6. They all share noble gas configurations.

Step-by-step solution

Determine Electron Configuration of Neutral Atoms

Before determining the electron configuration for the ions, identify the electron configurations for the neutral atoms F, P, Li, and Al. For neutral fluorine, the configuration is 1s2 2s2 2p5. For phosphorus, it's 1s2 2s2 2p6 3s2 3p3. For lithium, it's 1s2 2s1, and for aluminum, it's 1s2 2s2 2p6 3s2 3p1.

Adjust for Ionic Charge

For anions (negative ions), electrons are added to the electron configuration of the neutral atom. For cations (positive ions), electrons are removed. For F-, add one electron: 1s2 2s2 2p6. For P3-, add three electrons: 1s2 2s2 2p6 3s2 3p6. For Li+, remove one electron: 1s2. For Al3+, remove three electrons: 1s2 2s2 2p6.

Identify Commonality

The common feature in the electron configurations is that each ion has achieved a noble gas configuration. F-, with 1s2 2s2 2p6, mirrors the configuration of Ne. P3-, with 1s2 2s2 2p6 3s2 3p6, mimics the configuration of Ar. Li+, with 1s2, and Al3+, with 1s2 2s2 2p6, both have configurations identical to the noble gases He and Ne, respectively.

Key concepts

Ionic Charge Adjustment

Understanding how to adjust electron configurations for ions is crucial in the world of chemistry. When an atom gains or loses electrons, it becomes an ion, carrying a positive or negative charge. This process, known as ionic charge adjustment, is essential for grasping how ions achieve stability by reaching a noble gas configuration.

When an atom becomes a negatively charged ion, or anion, it has gained one or more electrons. These additional electrons fill up the available atomic orbitals. For example, the ion F^- includes an extra electron compared to its neutral state, resulting in the electron configuration of 1s2 2s2 2p6.

In contrast, a positively charged ion, or cation, forms when an atom loses electrons, which effectively reduces its electron configuration. For instance, the removal of an electron from a neutral lithium atom ( Li) leads to a Li⁺ ion, with an electron configuration of 1s2, resembling the helium (He) noble gas. The core concept here is the addition or subtraction of electrons to reflect the ion's charge, determining its stability and reactivity.

Noble Gas Configuration

Achieving a noble gas configuration is the holy grail for most atoms participating in chemical reactions, as it represents a state of maximum stability. Noble gases such as helium (He), neon (Ne), and argon (Ar) have complete outer electron shells, which makes them extraordinarily stable and largely nonreactive.

In the pursuit of stability, atoms will gain or lose electrons to emulate the electron configuration of the closest noble gas. This is evident in the provided exercise, where each ion adjusts its electron count to match that of a noble gas. For example, P^3- adds three electrons to achieve the same electron configuration as argon (Ar), 1s2 2s2 2p6 3s2 3p6.

This trend is a fundamental concept in chemistry because it is the driving force behind the formation of ionic bonds. When learning about electron configurations, recognizing an atom's desire to mimic noble gases facilitates a deeper understanding of ion formation and chemical bonding.

Electron Configuration Notation

The language used to describe the distribution of electrons in an atom's orbitals is called electron configuration notation. This system uses numbers and letters to specify the principal energy level (shell), subshell, and the number of electrons within that subshell.

The principal energy levels are denoted by numbers (1, 2, 3, ...), and subshells by lowercase letters (s, p, d, f). Each subshell can hold a fixed maximum number of electrons; for example, an 's' subshell can accommodate 2 electrons, while a 'p' subshell can hold 6. To illustrate, the notation 1s2 indicates the first energy level with a full 's' subshell, containing 2 electrons.

This notation is not only essential for writing out electron configurations of atoms and ions, but it is also important for understanding the underlying principles of chemical bonding and electronic transitions within an atom. Students should prioritize mastering this notation to effectively communicate and predict atomic behavior in the language of chemistry.