Question

Write the Lewis structure for ${\mathrm{O}}_2{\mathrm{F}}_2$. Predict the bond angles and hybridization of the two central oxygen atoms. Assign oxidation states and formal charges to the atoms in ${\mathrm{O}}_2{\mathrm{F}}_2$. The compound ${\mathrm{O}}_2{\mathrm{F}}_2$ is a vigorous and potent oxidizing and fluorinating agent. Are oxidation states or formal charges more useful in accounting for these properties of ${\mathrm{O}}_2{\mathrm{F}}_2$ ?

Short answer

The O2F2 molecule has a Lewis structure with each central oxygen atom forming a single bond with one fluorine atom and another oxygen atom. The central oxygen atoms exhibit sp³ hybridization and have bond angles less than 104.5° due to the repulsion from lone pairs. Central oxygen atoms have an oxidation state of +1 and a formal charge of +1, whereas fluorine atoms have an oxidation state of -1 and a formal charge of 0. The positive oxidation states of the central oxygen atoms are more useful for explaining the oxidizing and fluorinating properties of O2F2.

Step-by-step solution

Draw the Lewis Structure of O2F2

To draw the Lewis structure of O2F2, we first determine the total number of valence electrons present in the molecule. Each oxygen atom has 6 valence electrons, and each fluorine atom has 7 valence electrons. Thus, the total number of valence electrons in O2F2 is: Total valence electrons = (2 × 6) + (2 × 7) = 12 + 14 = 26 We'll place the oxygen atoms in the center of the molecule and the fluorine atoms on the outer sides. We can arrange 2 single bonds between the oxygen atoms and the fluorine atoms. Now, we start placing the lone pairs: - Each fluorine atom needs 3 lone pairs to complete its octet. - Each oxygen atom needs 2 lone pairs to complete its octet. The Lewis structure for O2F2 would look like this: F -- O -- O -- F | | O₂ O₂

Predict Bond Angles and Hybridization

To predict the bond angles and hybridization of the central oxygen atoms, we need to consider the electron domains surrounding each oxygen atom: - Each oxygen atom is bonded to two other atoms (the neighboring oxygen and one F atom). - Each oxygen atom also has 2 lone pairs. Therefore, each oxygen atom has 4 electron domains. According to VSEPR theory, an atom with 4 electron domains will have a tetrahedral electron geometry. However, each oxygen atom has one lone pair in its axial position which will push the bonding pairs down, giving it a bent molecular geometry. The ideal bond angle in a bent molecular geometry is approximately 104.5°, but the presence of the lone pairs could repulse the bonding electrons and slightly decrease the bond angle to a value below 104.5°. To find the hybridization, we observe that there are 4 electron domains around each oxygen atom. Thus, the oxygen atoms require three hybrid orbitals (one s and two p orbitals) to accommodate those domains. Therefore, the hybridization of the central oxygen atoms is sp³.

Assigning Oxidation States and Formal Charges

To find the oxidation state and formal charge for each atom, we follow these rules: 1) Oxidation state: - The oxidation state of fluorine is always -1. - For other oxygen atoms, the oxidation state is -2. We calculate the oxidation states: Central Oxygen: -2 (each) Fluorine: -1 (each) 2) Formal charge: - For oxygen atoms, the formal charge is the difference between the number of valence electrons (6) and the sum of the lone pair electrons and half the shared electrons. - For fluorine atoms, the formal charge is the difference between the number of valence electrons (7) and the sum of the lone pair electrons and half the shared electrons. We calculate the formal charges: Central Oxygen: $6-(2\times 2+\frac{1}{2}\times 2)=6-(4+1)=1$, Formal charge = +1 for each oxygen atom. Fluorine: $7-(3\times 2+\frac{1}{2}\times 2)=7-(6+1)=0$, Formal charge = 0 for each fluorine atom.

Understanding the Properties of O2F2

O2F2 is a strong oxidizing and fluorinating agent. The compound can break down to produce oxygen and fluorine gas, which are very reactive species. These reactive species can cause other substances to undergo oxidation and fluorination. The oxidation states of the oxygen atoms in O2F2 indicate that these atoms are less electronegative and more electropositive than their counterparts in other oxygen-containing compounds. This is because they carry a positive oxidation state (+1) rather than their typical oxidation state of -2. These central oxygen atoms with positive oxidation states help create a powerful oxidizing agent, making oxidation states more useful in accounting for the potent oxidizing and fluorinating properties of O2F2. In conclusion, oxidation states better account for the properties of O2F2, even though formal charges are important for understanding the distribution of charges within the molecule.