Question

Assign oxidation states to all of the atoms in each of the following: a. \(\mathrm{O}_{2}\) b. \(\mathrm{O}_{3}\) c. \(\mathrm{FeCl}_{3}\) d. \(\mathrm{FeCl}_{2}\)

Short answer

The oxidation states for the given compounds are as follows: a. \(\mathrm{O}_{2}\): Both oxygen atoms have an oxidation state of 0. b. \(\mathrm{O}_{3}\): Each oxygen atom has an average oxidation state of -2/3. c. \(\mathrm{FeCl}_{3}\): \(\mathrm{Fe}^{+3}\) and \(\mathrm{Cl}^{-1}\). d. \(\mathrm{FeCl}_{2}\): \(\mathrm{Fe}^{+2}\) and \(\mathrm{Cl}^{-1}\).

Step-by-step solution

Assign oxidation states to atoms in \(\mathrm{O}_{2}\)

In the case of a diatomic molecule such as \(\mathrm{O}_{2}\), each oxygen atom must have the same oxidation state. Since oxygen typically has an oxidation state of -2 and the overall charge of the molecule is 0, the oxidation states of each oxygen atom in \(\mathrm{O}_{2}\) must be 0 as they are in their elemental form.

Assign oxidation states to atoms in \(\mathrm{O}_{3}\)

In the ozone (\(\mathrm{O}_{3}\)) molecule, we notice that the structure is resonance-stabilized, meaning that electrons are delocalized among all the three oxygen atoms. As a result, the oxidation state of each oxygen atom is an average. Since the total charge of the molecule is 0, and we are dealing with three \(\mathrm{O}\) atoms, the average oxidation state for the oxygen atoms in ozone is (-2 / 3). Therefore, the oxidation states of the three oxygen atoms in \(\mathrm{O}_{3}\) are approximately: -2/3, -2/3, and -2/3. It is important to remember that this calculated oxidation state of -2/3 is an average and not always used in balancing redox reactions.

Assign oxidation states to atoms in \(\mathrm{FeCl}_{3}\)

In \(\mathrm{FeCl}_{3}\), we have an iron (Fe) atom and three chlorine (Cl) atoms. Since the overall charge of the molecule is 0, the oxidation state of iron plus three times the oxidation state of chlorine must equal zero. Chlorine typically has an oxidation state of -1. So the oxidation state of \(\mathrm{Fe}\) in this compound must be +3. Therefore, in \(\mathrm{FeCl}_{3}\) the oxidation states are \(\mathrm{Fe}^{+3}\) and \(\mathrm{Cl}^{-1}\).

Assign oxidation states to atoms in \(\mathrm{FeCl}_{2}\)

In \(\mathrm{FeCl}_{2}\), we have an iron (Fe) atom and two chlorine (Cl) atoms. Again, since the overall charge of the molecule is 0, the oxidation state of iron plus two times the oxidation state of chlorine must equal zero. Using the typical oxidation state of chlorine, which is -1, the oxidation state of iron in this compound must be +2. Consequently, in \(\mathrm{FeCl}_{2}\), the oxidation states are \(\mathrm{Fe}^{+2}\) and \(\mathrm{Cl}^{-1}\).

Key concepts

Chemical Compounds

Chemical compounds are substances formed when two or more different elements come together through chemical bonds. These bonds can be either ionic or covalent. For example, in

• \(\mathrm{FeCl}_{3}\) and \(\mathrm{FeCl}_{2}\), the compound formation involves the transfer of electrons due to the difference in electronegativity between iron (Fe) and chlorine (Cl). This results in ionic bonds.
• In contrast, \(\mathrm{O}_{2}\), a diatomic molecule, involves a covalent bond where two oxygen atoms share electrons.

Understanding how these atoms combine and the types of bonds they form is crucial to identifying oxidation states. Each atom or ion's charge contribution to a molecule plays a role in determining these states.

Redox Reactions

Redox reactions, short for reduction-oxidation reactions, are processes where there is a transfer of electrons between chemical species. These reactions involve two parallel processes:

• **Oxidation**, where a substance loses electrons and thus gains a more positive oxidation state.
• **Reduction**, where a substance gains electrons and thus decreases in oxidation state.

Assigning oxidation states to atoms is vital in tracking these reactions. For example:

• In \(\mathrm{FeCl}_{3}\), Cl is in its reduced form (\(-1\)), while Fe is oxidized to \(+3\).
• This means that Fe has lost electrons to Cl, illustrating a classic redox process.

By calculating oxidation states, we are able to understand and balance redox reactions, which is a fundamental aspect of chemistry.

Diatomic Molecules

Diatomic molecules consist of only two atoms, which can be the same element or different ones. The most common diatomic molecules in chemistry include the gases hydrogen (\(\mathrm{H}_{2}\), nitrogen (\(\mathrm{N}_{2}\), and oxygen (\(\mathrm{O}_{2}\).

• In the case of \(\mathrm{O}_{2}\), each oxygen atom has an oxidation state of zero because they share electrons equally, creating a stable bond with no overall charge.
• These molecules are essential in understanding basics of molecular chemistry as they provide a simple framework for studying more complex reactions.

In diatomic molecules involved in redox reactions, like \(\mathrm{O}_{2}\), the process of oxidation or reduction may alter the oxidation states from zero to another value as oxygen typically seeks to gain electrons.