Question

$$\begin{array}{rl}& \text{ Group the species that are isoelectronic: }{\mathrm{Be}}^{2+},{\mathrm{F}}^{-},{\mathrm{Fe}}^{2+}\\ & {\mathrm{N}}^{3-},\mathrm{He},{\mathrm{S}}^{2-},{\mathrm{Co}}^{3+},\mathrm{Ar}\end{array}$$

Short answer

The isoelectronic pairs are: ${\mathrm{Be}}^{2+}$ and $\mathrm{He}$; ${\mathrm{F}}^{-}$ and ${\mathrm{N}}^{3-}$; ${\mathrm{S}}^{2-}$ and $\mathrm{Ar}$; ${\mathrm{Fe}}^{2+}$ and ${\mathrm{Co}}^{3+}$.

Step-by-step solution

Define 'Isoelectronic'

Two species are said to be isoelectronic if they have the same number of electrons. To determine whether different species are isoelectronic, we need to count the total electrons present in each species.

Determine the Electron Count of Each Species

Calculate the number of electrons in each species: - ${\mathrm{Be}}^{2+}$: Beryllium has an atomic number of 4. After losing 2 electrons, it has 2 electrons.- ${\mathrm{F}}^{-}$: Fluorine has an atomic number of 9. After gaining 1 electron, it has 10 electrons.- ${\mathrm{Fe}}^{2+}$: Iron has an atomic number of 26. After losing 2 electrons, it has 24 electrons.- ${\mathrm{N}}^{3-}$: Nitrogen has an atomic number of 7. After gaining 3 electrons, it has 10 electrons.- $\mathrm{He}$: Helium naturally has 2 electrons.- ${\mathrm{S}}^{2-}$: Sulfur has an atomic number of 16. After gaining 2 electrons, it has 18 electrons.- ${\mathrm{Co}}^{3+}$: Cobalt has an atomic number of 27. After losing 3 electrons, it has 24 electrons.- $\mathrm{Ar}$: Argon naturally has 18 electrons.

Group Isoelectronic Species

Identify species with the same electron count:- 2 electrons: ${\mathrm{Be}}^{2+}$ and $\mathrm{He}$- 10 electrons: ${\mathrm{F}}^{-}$ and ${\mathrm{N}}^{3-}$- 18 electrons: ${\mathrm{S}}^{2-}$ and $\mathrm{Ar}$- 24 electrons: ${\mathrm{Fe}}^{2+}$ and ${\mathrm{Co}}^{3+}$

Key concepts

Electron Count

The concept of electron count is crucial when discussing isoelectronic species. Isoelectronic species share the same number of electrons, resulting in similar electron configurations even though they may belong to different elements or ions. For example:

• Beryllium (${\mathrm{Be}}^{2+}$) has an atomic number of 4, therefore it typically has 4 electrons. When it loses 2 electrons, it ends up with 2 electrons.
• Sulfur (${\mathrm{S}}^{2-}$) has 16 electrons in its neutral state, but gains two additional electrons, reaching a total of 18 electrons.

Understanding how to count electrons not only helps in identifying isoelectronic species, but also in predicting chemical behavior and bonding characteristics.

Atomic Number

The atomic number of an element is fundamental to understanding its properties because it denotes the number of protons in the nucleus of an atom. This is inherently equal to the number of electrons in a neutral atom, which determines its electron configuration.

• Fluorine (${\mathrm{F}}^{-}$)
has an atomic number of 9. As a negatively charged ion (${\mathrm{F}}^{-}$), it gains an extra electron, totaling 10 electrons.
• Argon ($\mathrm{Ar}$)
has an atomic number of 18, meaning it has 18 protons and maintains 18 electrons in its neutral, stable state.

The atomic number is a key factor when assessing electron configurations and chemical identity.

Ionic Charge

Ionic charge arises from an atom gaining or losing electrons, leading to a net positive or negative charge.

• Cations, like Iron (${\mathrm{Fe}}^{2+}$)
, have lost electrons, resulting in a positive charge. Here, iron loses 2 electrons, reducing from 26 electrons to 24.
• Anions, such as Nitrogen (${\mathrm{N}}^{3-}$)
, have gained electrons, leading to a negative charge. Nitrogen gains 3 electrons, increasing from 7 to 10 electrons.

The ionic charge is essential not only for determining if species are isoelectronic but also influences chemical reactivity and stability.

Electron Configuration

Electron configuration provides a detailed map of where each electron resides within an atom or ion. While the electron count is crucial, the distribution across energy levels (shells) gives more insight into chemical behavior. For instance:

• Argon's ($\mathrm{Ar}$) electron configuration
is $1s^{2}2s^{2}2p^{6}3s^{2}3p^6$, reflecting its noble gas stability.
• Sulfur's (${\mathrm{S}}^{2-}$) electron configuration
after gaining 2 electrons becomes $1s^{2}2s^{2}2p^{6}3s^{2}3p^6$, mimicking argon due to its isoelectronic nature.

The electron configuration not only confirms isoelectronic species but also aids in predicting bonds and molecular structures.