Question

Lewis theory is satisfactory to explain bonding in the ionic compound \(\mathrm{K}_{2} \mathrm{O},\) but it does not readily explain formation of the ionic compounds potassium superoxide, \(\mathrm{KO}_{2}\), and potassium peroxide, \(\mathrm{K}_{2} \mathrm{O}_{2}\) (a) Show that molecular orbital theory can provide this explanation. (b) Write Lewis structures consistent with the molecular orbital explanation.

Short answer

Molecular Orbital theory explains the formation of KO2 and K2O2 by considering that Potassium atoms donate their 4s electrons to the antibonding orbitals of the Oxygen molecules, forming ionic compounds. The Lewis structures confirm these ionic bonds, showing electrons' transfer from Potassium to Oxygen molecules.

Step-by-step solution

Understanding molecular orbital theory for KO2

For KO2, Potassium (K) has one valence electron in its 4s orbital. O2 (dioxygen) is a diradical and exhibits paramagnetism. It has corresponding π* (antibonding) orbitals populated by one electron each. When Potassium reacts with O2, the electron from 4s orbital of K goes to one of these antibonding orbitals, making O2 a full bonding species while giving K a +1 charge, forming an ionic compound, KO2.

Understanding molecular orbital theory for K2O2

For K2O2, each Potassium atom donates one electron to an Oxygen molecule. As a result, each K atom gets a +1 charge, and O2 gets a -1 charge. Since two K atoms donate one electron each, the overall charge forms as neutral, creating KO2.

Writing Lewis structures for KO2

For KO2, the Lewis structure will show that Oxygen (O2) molecule forms an ionic bond with Potassium (K+), represented by placing K+(s) Naear [O2]-. It indicates that one electron is transferred from Potassium to Oxygen molecule.

Writing Lewis structures for K2O2

For K2O2, the Lewis structure will show two Potassium ions (K+) forming an ionic bond with Peroxide (O22-), represented by placing two K+ near [O - O]2-. It indicates that an electron from each Potassium atom is transferred to the Oxygen molecule.

Key concepts

Lewis Structures

Lewis Structures are a simple and effective way to understand the covalent bonds in a molecule. This method uses dots and lines to represent electrons and bonds between atoms. Each dot corresponds to a valence electron, while a line represents a shared pair of electrons forming a covalent bond.

To construct a Lewis structure for any compound:

• List all the valence electrons available for bonding in the atoms involved.
• Connect atoms primarily through single bonds, which are represented as lines.
• Ensure that the octet rule is satisfied; each atom should typically have eight electrons surrounding it, except for hydrogen, which should have two.
• Use dots around individual atoms to represent electrons that are not shared.

For example, in \( \text{KO}_2 \), the oxygen molecule (\( \text{O}_2 \)) shares electrons with potassium (\( \text{K}^+ \)), causing the oxygen atoms to gain negative charges as one of their antibonding orbitals is filled.

Ionic Compounds

Ionic compounds form from the electrostatic attraction between cations and anions. This typically occurs when electrons are transferred from a metal to a non-metal atom.

Metal atoms, like potassium (K), tend to lose electrons to form positively charged ions or cations. Non-metals, such as oxygen, gain these electrons to form negatively charged ions or anions. The charge difference creates a strong attraction that holds the compound together.

In \( \text{K}_2\text{O}_2 \), potassium ions (\( \text{K}^+ \)) bond with peroxide ions (\( \text{O}_2^{2-} \)), forming a stable ionic compound. Each potassium ion loses one electron, leading to the ionic bond formation with the peroxide ion, which stabilizes with the gain of electrons.

Bonding in Chemistry

Bonding in chemistry refers to the interactions that hold atoms together in molecules. There are several types of chemical bonds: ionic, covalent, and metallic bonds.

In ionic bonding, as mentioned earlier, electrons transfer between atoms, creating ions that attract each other. Covalent bonds involve the sharing of electrons. Metallic bonds, on the other hand, occur between atoms of metallic elements, where electrons are free to move around.

The nature of a bond can determine a compound’s physical and chemical properties, such as melting point, solubility, and reactivity. In compounds like \( \text{KO}_2 \) and \( \text{K}_2\text{O}_2 \), both ionic and covalent characters play roles, particularly in how oxygen handles extra electrons from potassium, enhancing their stability.

Molecular Orbital Explanation

The Molecular Orbital (MO) theory provides a more detailed understanding of bonding, especially in complex molecules that exhibit both ionic and covalent character.

In MO theory, atomic orbitals combine to form molecular orbitals that are spread over the entire molecule, accommodating the bonded electrons. Electrons fill these orbitals starting from the lowest energy level.

For example, in \( \text{KO}_2 \), the electron from potassium's 4s orbital moves to one of the antibonding π* orbitals of the \( \text{O}_2 \) molecule, stabilizing the \( \text{O}_2 \) ion with a single negative charge and giving potassium a positive charge. This electron movement creates the ionic characteristic of the compound, yet oxygen preserves its covalent bond within the ion.

Molecular Orbitals’ explanations simplify why some compounds, like potassium superoxide and potassium peroxide, show unique bonding patterns that aren't easily explained using just Lewis Structures.