Problem 133 - Chemistry: An Atoms-Focused Approach

Question

In the following species, which atom is most likely to have an unpaired electron? (a) \(\mathrm{SO}^{+} ;\) (b) \(\mathrm{NO} ;\) (c) \(\mathrm{CN} ;\) (d) \(\mathrm{OH}.\)

Step-by-Step Solution

Answer

Answer: Sulfur in the SO+ species is most likely to have an unpaired electron.

Step 1: Electron Configurations of Atoms in their Ground State

Write down the electron configurations of S, O, N, C, and H in their ground states. S: [Ne]\(3s^2$$3p^4\) O: [He]\(2s^2$$2p^4\) N: [He]\(2s^2$$2p^3\) C: [He]\(2s^2$$2p^2\) H: \(1s^1\)

Step 2: Determine Electron Configurations in Each Species

Determine the electron configurations of the atoms in the given species. (a) SO\(^+\): S lost one electron and O gained one electron, making their configurations: S: [Ne]\(3s^2$$3p^3\) O: [He]\(2s^2$$2p^5\) (b) NO: N gained one electron and O lost one electron, making their configurations: N: [He]\(2s^2$$2p^4\) O: [He]\(2s^2$$2p^3\) (c) CN: C triple bonded to N implies that C shares 3 pairs of electrons, making their configurations: C: [He]\(2s^2$$2p^4\) (sharing 3 electron pairs) N: [He]\(2s^2$$2p^6\) (sharing 3 electron pairs) (d) OH: O and H are bonded with a single bond, making their configurations: O: [He]\(2s^2$$2p^5\) (sharing 1 electron pair) H: \(1s^2\) (sharing 1 electron pair)

Step 3: Identify Atoms with Unpaired Electrons

Look for atoms that have an unpaired electron. (a) SO\(^+\): S: [Ne]\(3s^2$$3p^3\) (unpaired) O: [He]\(2s^2$$2p^5\) (paired) (b) NO: N: [He]\(2s^2$$2p^4\) (paired) O: [He]\(2s^2$$2p^3\) (unpaired) (c) CN: C: [He]\(2s^2$$2p^4\) (paired) N: [He]\(2s^2$$2p^6\) (paired) (d) OH: O: [He]\(2s^2$$2p^5\) (paired) H: \(1s^2\) (paired)

Step 4: Identify the Atom most likely to have an Unpaired Electron

Find the atom that has the highest probability of having an unpaired electron. From the previous step, only S in SO\(^+\) and O in NO have unpaired electrons. Considering other factors such as electronegativity and ionization energy, oxygen has a higher electronegativity than sulfur, making it more likely to have a paired electron. Therefore, the atom most likely to have an unpaired electron is Sulfur in the SO\(^+\) species.

Key Concepts

Electron Configuration

Electron configuration is a way to represent where electrons reside in an atom. It helps us understand how electrons are distributed among the various atomic orbitals. Each electron in an atom is organized into shells and subshells, which include the s, p, d, and f types.

The notation usually starts with the number of the outermost shell, followed by the letter representing the orbital type, and last is a superscript denoting the number of electrons in that orbital.

For instance, Oxygen has the electron configuration \(\text{[He]} 2s^2 2p^4\). This notation tells us about the energy level (the number preceding s or p), the type of orbital (s or p), and how many electrons are in that orbital.

Understanding electron configuration is key to predicting how atoms will bond with each other and helps in identifying which atoms might have unpaired electrons.

Molecular Species

In chemistry, molecular species refer to groups of atoms bonded together and functioning as a single entity. Each molecular species involves atoms sharing, gaining, or losing electrons, leading to different configurations compared to isolated atoms.

For example, in the species SO\(^+\), the sulfur atom loses an electron, which changes its electron configuration compared to a neutral sulfur atom. This can lead to the presence of unpaired electrons.

When atoms bond within a species, they form molecules, and the stability of these molecules depends on their electron configurations and how the electrons are shared or distributed.

Molecular species incorporate changes from the individual electron configurations of atoms in their isolated states, which can further influence chemical reactivity and physical properties.

Ground State Atoms

Ground state atoms are atoms that have their electrons in the lowest possible energy levels. This is the most stable configuration because there is the least energy in the system.

The ground state configuration aligns with the rules of filling orbitals: electrons prefer to go in the lowest energy orbital available and tend to have opposite spins when paired.

For instance, the ground state configuration of nitrogen is \[\text{[He]} 2s^2 2p^3\], which signifies that the 2p orbital contains all unpaired electrons.

When considering chemical reactions or bonding with other atoms, atoms often initially engage from their ground states. However, bonding can change their electron configurations either by pairing up unpaired electrons or by altering distribution through ionic or covalent bonds.

Bonding and Antibonding Orbitals

The concepts of bonding and antibonding orbitals are essential in molecular orbital theory, a method for determining molecular structure.

Bonding orbitals are lower energy orbitals formed when atomic orbitals combine constructively, leading to more stability and promoting the bond between atoms.

Antibonding orbitals are higher in energy, resulting from the destructive combination of atomic orbitals, which can weaken or prevent a bond.

Electrons that reside in bonding orbitals help to hold atoms together, contributing to the stability of the molecule. In contrast, electrons in antibonding orbitals can destabilize the molecule. Understanding the distribution of electrons between these orbitals can explain molecular properties such as magnetism and reactivity.