Question

Write the Lewis structure for \(\mathrm{O}_{2} \mathrm{~F}_{2}\left(\mathrm{O}_{2} \mathrm{~F}_{2}\right.\) exists as \(\mathrm{F}-\mathrm{O}-\mathrm{O}-\mathrm{F}\) ). Assign oxidation states and formal charges to the atoms in \(\mathrm{O}_{2} \mathrm{~F}_{2}\). This compound 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 Lewis structure for O2F2 is \[F-O-O-F\]. The oxidation states are F(-1)-O(-1)-O(-1)-F(-1) and the formal charges are F(0)-O(0)-O(0)-F(0). The oxidation states are more useful in accounting for O2F2's properties as a potent oxidizing and fluorinating agent due to the unusual peroxide state of the oxygen atoms with an oxidation state of -1, which makes them more reactive.

Step-by-step solution

Draw the Lewis structure for O2F2

Since the molecule is given as F-O-O-F, the first step is to draw the Lewis structure for O2F2. Oxygen has 6 valence electrons and Fluorine has 7 valence electrons. In total, there are 26 valence electrons in this molecule. We will place the electrons as follows: 1. Connect the two oxygen atoms with a single bond, which will use 2 electrons. 2. Connect each fluorine atom to its respective oxygen atom with a single bond, using 4 more electrons (2 for each bond). 3. Distribute the remaining 20 electrons around the atoms as lone pairs. Each fluorine atom will have 6 non-bonding electrons (3 lone pairs), and each oxygen atom will have 4 non-bonding electrons (2 lone pairs); 8 electrons will be placed around each oxygen atom. Lewis structure for O2F2: \[F-O-O-F\] Here, each oxygen has 2 lone pairs and each fluorine has 3 lone pairs.

Assign oxidation states to each atom

To calculate the oxidation states for each atom in the O2F2 molecule, we can use the following rules: 1. The oxidation state of an element in its free state is 0. 2. The oxidation state of a monatomic ion is equal to its charge. 3. Oxygen is usually assigned an oxidation state of -2 (except peroxides where it is -1 and in compounds with F where it can be positive). 4. Fluorine is always assigned an oxidation state of -1. Here, we have two peroxide oxygen atoms (oxidation state of -1 each), and two fluorine atoms (oxidation state of -1 each). Oxidation states: F(-1)-O(-1)-O(-1)-F(-1)

Assign formal charges to each atom

To calculate the formal charges of each atom, we use the formula: Formal charge = Valence electrons - (0.5 × Bonding electrons + Lone pair electrons) For F1: Valence electrons = 7 (F has 7 valence electrons) Bonding electrons = 2 (1 bond with O) Lone pair electrons = 6 (3 lone pairs) Formal charge F1 = 7 - (0.5 × 2 + 6) = 0 For O1: Valence electrons = 6 (O has 6 valence electrons) Bonding electrons = 4 (1 bond with F1, 1 bond with O2) Lone pair electrons = 4 (2 lone pairs) Formal charge O1 = 6 - (0.5 × 4 + 4) = 0 For O2 : Formal charge O2 = 6 - (0.5 × 4 + 4) = 0 (same as for O1) For F2 : Formal charge F2 = 7 - (0.5 × 2 + 6) = 0 (same as for F1) Formal charges: F(0)-O(0)-O(0)-F(0)

Compare oxidation states and formal charges to explain O2F2's properties

The oxidation states and formal charges for O2F2 are as follows: Oxidation states: F(-1)-O(-1)-O(-1)-F(-1) Formal charges: F(0)-O(0)-O(0)-F(0) While the formal charges indicate all atoms have a neutral charge, the oxidation states show that all atoms have negative oxidation states. In this case, the oxidation states are more useful in accounting for O2F2's properties as a potent oxidizing and fluorinating agent. This is because it is easier to see that the oxygen atoms are in an unusual state with an oxidation state of -1 (peroxide), which makes them more reactive. This reactivity results in the strong oxidizing and fluorinating properties of the O2F2 molecule.

Key concepts

Oxidation States

Understanding oxidation states is critical when analyzing the chemical behavior of compounds such as dioxygen difluoride (O2F2). Oxidation states, often referred to as oxidation numbers, provide insight into the electron distribution in molecules and the degree of oxidation or reduction of atoms within the molecule.

For O2F2, determining oxidation states involves following general rules: fluorine, with the highest electronegativity, is assigned an oxidation state of -1 since it typically gains an electron when forming compounds. Conversely, oxygen usually has an oxidation state of -2, but in the case of peroxides like O2F2, it adopts an oxidation state of -1 due to the O-O bond. These assignments are essential in understanding the reactivity of the molecule, especially its tendency to act as an oxidizing agent.

• Oxygen (peroxide linkage): -1
• Fluorine: -1

Once we've established the oxidation states, we should note that both oxygen atoms in peroxide linkage have an oxidation state of -1, which is higher than the usual -2. This less stable configuration indicates a propensity to attain a more stable state by accepting electrons, thereby driving the compound's oxidizing behavior.

Formal Charge

The formal charge of an atom in a molecule helps us predict the stability of the molecule, possible arrangements of atoms, and the reactivity.

The formal charge is calculated by the formula: Formal charge = Valence electrons - (0.5 × Bonding electrons + Lone pair electrons). For O2F2, each fluorine and oxygen atom has a formal charge of zero, which may initially suggest a stable configuration.

However, formal charges are most informative when considered in conjunction with oxidation states. Despite all atoms having a formal charge of zero, indicating no strong electrical imbalances within the molecule, this doesn't account for the oxidizing nature of O2F2. It is the oxidation states that provide a clearer picture of that particular property.

Potent Oxidizing Agent

The property of being a potent oxidizing agent relates directly to a substance's ability to accept electrons during chemical reactions, thereby oxidizing another species. Dioxygen difluoride (O2F2) is known for its vigorous and potent oxidizing power.

Looking at O2F2's structure, we understand that the unusual oxidation state of oxygen at -1 (due to the presence of the peroxide bond) makes it electron-deficient and therefore reactive. This reactivity fuels its capability to act as an oxidizing agent, as these oxygen atoms look to 'gain' electrons to revert to a more stable -2 state.

• The peroxide bond in O2F2 is a source of instability and reactivity.
• Its oxidizing ability is pronounced owing to the oxygen atoms' higher oxidation states.

When assessing the chemical behavior of compounds like O2F2, the oxidation states give more detailed insight into their reactive nature, particularly as potent oxidizing agents, than formal charges. This underscores the molecule's potential to engage in highly energetic reactions, making it a powerful tool for various chemical applications.