Problem 135 - Introductory Chemistry Atoms First

Question

The oxide ion is \(\mathrm{O}^{2-}\). How does this ion differ from the peroxide ion, \(\mathrm{O}_{2}^{2-} ?\) Draw dot diagrams for both.

Step-by-Step Solution

Answer

The peroxide ion (O₂²⁻) differs from the oxide ion (O²⁻) in that it consists of two oxygen atoms linked by a single bond, with each oxygen atom carrying a partial negative charge. The oxide ion has a full octet of electrons around the oxygen atom, while in the peroxide ion, each oxygen atom shares two electrons for a single bond in addition to having six nonbonding valence electrons. The Lewis dot diagrams for the oxide ion and the peroxide ion are: \( [\underset{.}{\overset{.}{O}}\: :]^{2-} \) and \( [ \underset{.}{\overset{.}{O}} - \underset{.}{\overset{.}{O}} :]^{2-} \), respectively.

Step 1: Electron Configuration for \(\mathrm{O}^{2-}\) and \(\mathrm{O}_{2}^{2-}\)

The electron configuration for a neutral oxygen atom is 1s² 2s² 2p⁴. To form the oxide ion (O²⁻), it gains two electrons to fill its outer shell, giving a total of 10 electrons. So, the electron configuration for \(\mathrm{O}^{2-}\) is 1s² 2s² 2p⁶. On the other hand, the peroxide ion is formed by two oxygen atoms sharing two electrons to create a single bond between them (O-O). Each oxygen atom also gains one more electron, which brings the total number of electrons in the peroxide ion to 12 (6 for each oxygen atom). The electron configuration for \(\mathrm{O}_{2}^{2-}\) is the same as that of \(\mathrm{O}^{2-}\) for each oxygen atom, which is 1s² 2s² 2p⁶.

Step 2: Drawing Lewis Dot Diagrams for \(\mathrm{O}^{2-}\) and \(\mathrm{O}_{2}^{2-}\)

In this step, we'll draw the Lewis dot diagrams for both ions. For the oxide ion: 1. Write the oxygen symbol (O) and surround it with eight dots, representing the eight valence electrons. 2. Since it's a negatively charged ion, add brackets around the ion and put the charge on the upper right corner. The Lewis dot diagram for the oxide ion (O²⁻) will look like this: \[ [\underset{.}{\overset{.}{O}}\: :]^{2-} \] For the peroxide ion: 1. Write two oxygen symbols (O) with two dots between them, representing the single bond formed by the shared electrons. 2. Add six more dots around each oxygen atom, making sure each oxygen atom has eight electrons in total. 3. As before, add brackets around the ion and put the charge on the upper right corner. The Lewis dot diagram for the peroxide ion (O₂²⁻) will look like this: \[ [ \underset{.}{\overset{.}{O}} - \underset{.}{\overset{.}{O}} :]^{2-} \]

Step 3: Comparing the Ions and Identifying Differences

Now that we have the Lewis dot diagrams for both ions, we can compare them and identify their differences: 1. The oxide ion (O²⁻) consists of a single oxygen atom with a charge of 2-, while the peroxide ion (O₂²⁻) consists of two oxygen atoms with a total charge of 2-. 2. The oxide ion has a full octet of electrons around the oxygen atom, while in the peroxide ion, each oxygen atom shares two electrons for a single bond in addition to having six nonbonding valence electrons. In conclusion, the peroxide ion (O₂²⁻) differs from the oxide ion (O²⁻) in having two oxygen atoms linked by a single bond, with each oxygen atom carrying a partial negative charge.

Key Concepts

Electron Configuration

Electron configuration is an essential concept in understanding how atoms and ions form. It describes the arrangement of electrons in an atom's orbitals. For a neutral oxygen atom, the electron configuration is given by 1s² 2s² 2p⁴. This indicates the distribution of oxygen's eight electrons in its atomic orbitals.
To form the oxide ion, \(\mathrm{O}^{2-}\), oxygen gains two additional electrons to achieve a stable octet. Consequently, the electron configuration becomes 1s² 2s² 2p⁶, totaling ten electrons. This configuration results in a full outer shell, which is characteristic of a stable ion.
In contrast, the peroxide ion, \(\mathrm{O}_{2}^{2-}\), involves two oxygen atoms. Each oxygen atom retains its full octet by sharing two electrons to form a bond (O-O). The electron configuration for each oxygen in this peroxide ion is similarly 1s² 2s² 2p⁶, resulting in twelve total electrons for the ion. Thus, while the electron configuration for the oxide and peroxide ions is the same per oxygen when in these ions, the structural arrangement of these atoms makes them distinct.

Lewis Dot Diagrams

Lewis Dot Diagrams are a useful tool for visualizing the valence electrons in atoms and molecules. They help in understanding how electrons are shared or transferred between atoms to form chemical bonds.
For the oxide ion, \(\mathrm{O}^{2-}\), the Lewis dot diagram is constructed by placing the oxygen symbol (O) and surrounding it with eight dots, corresponding to its valence electrons. Since it is a negatively charged ion, place brackets around the dot diagram indicating that the ion carries a \(2^{-}\) charge. This confirms a complete octet around the oxygen atom.
On the other hand, the Lewis dot diagram for the peroxide ion, \(\mathrm{O}_{2}^{2-}\), involves two oxygen (O) atoms. These atoms share a pair of electrons, forming a single bond depicted by two dots or a line between them. Each oxygen atom is surrounded by six more dots, ensuring both atoms achieve a complete octet. The entire ion is set within brackets to denote the \(2^{-}\) charge. Understanding these representations highlights the differences in bonding and electron sharing between oxide and peroxide ions.

Oxygen Ion Bonding

Oxygen ion bonding is a key concept in understanding different oxygen compounds. In ions like oxide and peroxide, the bond formation involves electron transfer and sharing that stabilize these species.
The oxide ion, \(\mathrm{O}^{2-}\), features a straightforward ionic bond. Oxygen gains two electrons, achieving a stable electron arrangement. This gain gives the oxide ion its \(2^{-}\) charge, characterizing it by a complete outer electron shell.
Conversely, in the peroxide ion, \(\mathrm{O}_{2}^{2-}\), the bond formation involves both atoms sharing electrons to form a covalent bond. The electrons shared between the two oxygen atoms create a single bond, while each atom has six nonbonding (or lone) electrons. These characteristics lead to partial negative charges on each oxygen atom, differing from the full negative charge found in the oxide ion.
These forms of bonding underscore the versatility of oxygen in forming varied compounds by gaining, losing, or sharing electrons.