Question

The neutral OH molecule has been implicated in certain ozonedestroying processes that take place in the upper atmosphere. (a) Draw electron-dot structures for the OH molecule and the \(\mathrm{OH}^{-}\) ion. (b) Electron affinity can be defined for molecules just as it is defined for single atoms. Assuming that the electron added to \(\mathrm{OH}\) is localized in a single atomic orbital on one atom, identify which atom is accepting the electron, and give the \(n\) and \(l\) quantum numbers of the atomic orbital. (c) The electron affinity of \(\mathrm{OH}\) is similar to but slightly more negative than that of \(\mathrm{O}\) atoms. Explain.

Short answer

O atom receives electron (2p orbital, \(n=2\), \(l=1\)); OH's electron affinity is more negative due to bond stabilization.

Step-by-step solution

Draw the Lewis Structure of OH

For the OH molecule, the oxygen atom has 6 valence electrons and hydrogen has 1 valence electron. When forming a bond, they share one electron pair. Draw the oxygen atom with its valence electrons, sharing a pair of electrons with the hydrogen's single electron, forming an O-H bond. This results in oxygen having 6 total electrons including the shared pair.

Draw the Lewis Structure of OH⁻

For the OH⁻ ion, incorporate the extra electron into the structure. Place the additional electron on the oxygen atom, forming a negative charge on the oxygen, giving it a total of 7 electrons with one pair being shared with hydrogen to form the bond.

Identify Atom Receiving the Extra Electron

In the OH molecule, adding an electron to form OH⁻, the extra electron is localized on the oxygen atom. This is because oxygen is more electronegative and can stabilize the extra electron better than hydrogen.

Determine Quantum Numbers for the Electron

The additional electron in OH⁻ is added to an oxygen orbital. The oxygen atom has its valence electrons in the 2p orbitals; hence, the added electron will likely go into a 2p orbital. Therefore, the quantum numbers are: principal quantum number, \(n = 2\), and angular momentum quantum number, \(l = 1\).

Explain the Electron Affinity of OH

The electron affinity of OH is slightly more negative than that of an isolated O atom because the oxygen in OH can better stabilize the additional electron through the polar covalent bond with hydrogen. This partial stabilization effectively makes it energetically more favorable to add an electron to OH.

Key concepts

Electron-Dot Structure

The electron-dot structure, also known as a Lewis structure, is a simple way to represent the valence electrons of atoms in a molecule. It shows which electrons are shared between atoms and which electrons remain unshared (lone pairs). When drawing these structures, it's important to consider the valence electrons, as they are the ones involved in bond formation.

For instance, in the case of the OH molecule, oxygen has six valence electrons, while hydrogen has just one. In order to form a bond, they share a pair of electrons. Thus, the Lewis structure for OH includes a single bond between oxygen and hydrogen and four unshared electrons represented as dots around the oxygen. For the hydroxide ion (OH⁻), we add an extra electron to the oxygen, creating a full valence shell and imparting a negative charge. This additional electron creates an extra lone pair on oxygen, an essential element that influences the chemical behavior of the ion.

Electron Affinity

Electron affinity refers to the amount of energy released when an atom or molecule gains an electron. In the context of a neutral molecule like OH, electron affinity can be an important measure. When OH gains an electron, it forms the hydroxide ion (OH⁻). The extra electron is added to the oxygen atom, since oxygen, being more electronegative than hydrogen, is better suited to accommodate the electron.

Electron affinity in molecules can be compared to that in individual atoms. In this scenario, the electron affinity of OH is slightly more negative than that of an isolated oxygen atom. This means that adding an electron to OH releases a bit more energy than adding it to a single oxygen atom. The reason lies in the polar covalent bond between oxygen and hydrogen, which helps to stabilize the additional electron through partial charge separation. Thus, molecules with polar covalent bonds often exhibit different electron affinity behavior compared to single atoms.

Quantum Numbers

Quantum numbers are critical in describing the position and energy of an electron within an atom. There are four quantum numbers, but only two are directly involved in this context: the principal quantum number ( ) and the angular momentum quantum number ( l).
The principal quantum number ( ) indicates the energy level or shell where the electron resides. For the valence electrons of oxygen, this is 2, since they are located in the second shell. The angular momentum quantum number ( l) specifies the subshell or orbital type, with values ranging from 0 to n-1. In this case, the electron is added to a p orbital, giving an l value of 1.

When the OH molecule gains an electron to become OH⁻, this electron is placed in one of the 2p orbitals of oxygen. Hence, the quantum numbers for the additional electron are = 2 and l = 1, perfectly positioning the electron within the established second shell, in one of the three possible p orbitals.

Polarity in Molecules

Polarity in molecules is influenced by the distribution of electronic charges within the molecule. A molecule is polar if it has a positive charge on one end and a negative charge on the other, resulting in a dipole moment. This is typically due to differences in electronegativity between the bonding atoms.
In the OH molecule, oxygen is significantly more electronegative than hydrogen, which results in a partial negative charge (δ⁻) on the oxygen atom and a partial positive charge (δ⁺) on the hydrogen atom. This uneven electron sharing creates a dipole moment, making the molecule polar.

Polarity is important because it affects how molecules interact with each other, such as when they dissolve in water or interact in biochemical reactions. Polar molecules tend to dissolve well in water due to their ability to form hydrogen bonds and other polar interactions, which is essential for various biological and chemical processes.