Question

(a) Determine the formal charge on the chlorine atom in the hypochlorite ion, \(\mathrm{ClO}^{-},\) and the perchlorate ion, \(\mathrm{ClO}_{4}^{-},\) using resonance structures where the \(\mathrm{Cl}\) atom has an octet. (b) What are the oxidation numbers of chlorine in \(\mathrm{ClO}^{-}\) and in \(\mathrm{ClO}_{4}^{-} ?\) (c) Is it uncommon for the formal charge and the oxidation state to be different? Explain. (d) Perchlorate is a much stronger oxidizing agent than hypochlorite. Would you expect there to be any relationship between the oxidizing power of the oxyanion and either the oxidation state or the formal charge of chlorine?

Short answer

In the hypochlorite ion, \(\mathrm{ClO}^{-}\), the formal charge on chlorine is 0, and the oxidation number is +1. In the perchlorate ion, \(\mathrm{ClO}_{4}^{-}\), the formal charge on chlorine is 3, and the oxidation number is +7. It is not uncommon for the formal charge and oxidation state to be different, as they represent different properties. The oxidizing power of oxyanions is related to the oxidation state rather than the formal charge of chlorine. In this case, perchlorate is a stronger oxidizing agent due to its higher oxidation state.

Step-by-step solution

Determine the formal charge of chlorine in hypochlorite ion and perchlorate ion

To find the formal charge of an atom in a molecule or ion, use the formula: Formal Charge = (# of valence electrons in free atom) - (# of lone pair electrons) - (1/2 * # of bonding electrons). We also need to consider the resonance structures for both ions where chlorine has an octet. (a) For hypochlorite ion, \(\mathrm{ClO}^{-}\): The Lewis structure has single bond between chlorine and oxygen atoms, and chlorine has three lone pairs of electrons while oxygen has two lone pairs and one extra electron due to its negative charge. 1. Formal charge of Cl: \(7 - (6 + (1/2 * 2)) = 7 - 7 = 0\) 2. Formal charge of O: \(6 - (4 + (1/2 * 2)) = 6 - 5 = 1\) For perchlorate ion, \(\mathrm{ClO}_{4}^{-}\): The Lewis structure has single bonds between chlorine and four oxygen atoms, and each oxygen has two lone pairs of electrons with one extra electron due to the negative charge. Chlorine has no lone pairs of electrons in this ion. 1. Formal charge of Cl: \(7 - (0 + (1/2 * 8)) = 7 - 4 = 3\) 2. Formal charge of each O: \(6 - (4 + (1/2 * 2)) = 6 - 5 = 1\)

Calculate the oxidation number of chlorine in hypochlorite ion and perchlorate ion

To find the oxidation number, assign oxidation numbers for all atoms in the compound or ion. The sum of the oxidation numbers is equal to the charge of the compound or ion. (b) Oxidation number of Cl in \(\mathrm{ClO}^{-}\): Let the oxidation number of Cl be x. Since the ion carries a negative charge, x + (-2) = -1 --> x = +1. So, the oxidation number of Cl in \(\mathrm{ClO}^{-}\) is +1. Oxidation number of Cl in \(\mathrm{ClO}_{4}^{-}\): Let the oxidation number of Cl be x. The ion carries a negative charge, x + 4(-2) = -1 --> x = +7. So, the oxidation number of Cl in \(\mathrm{ClO}_{4}^{-}\) is +7.

Discuss the difference between formal charge and oxidation state

(c) It is not uncommon for the formal charge and the oxidation state to be different. Formal charge is the difference between the number of valence electrons in a free atom and the number of electrons assigned to an atom in a molecule or ion. Oxidation state is an assigned number which helps to determine the degree of oxidation or reduction of an atom in a chemical species.

Analyze the relationship between oxidizing power and formal charge or oxidation state of chlorine

(d) Oxidizing power is the ability of a chemical species to accept electrons or oxidize another species. A higher oxidation state of an element means a higher potential to be reduced (gain electrons), which indicates a stronger oxidizing power. However, formal charge is not directly related to the oxidizing power of the ion. In this case, the perchlorate ion has a higher oxidation state (+7) than the hypochlorite ion (+1), making it a stronger oxidizing agent. This shows that the oxidizing power of oxyanion is related to the oxidation state, rather than the formal charge of chlorine.

Key concepts

Oxidation Number

The oxidation number of an atom in a molecule or ion represents the number of electrons that an atom gains or loses during a chemical reaction. It's a bookkeeping method used to track electron transfers. To determine oxidation numbers, we typically assume the oxidation number for oxygen is -2 unless stated otherwise. Accordingly:

• In \( \mathrm{ClO}^{-} \), the oxidation number of chlorine is calculated by setting up the equation: \( x + (-2) = -1 \). Solving for \( x \), the oxidation number is +1.


• For \( \mathrm{ClO}_{4}^{-} \), the equation is \( x + 4(-2) = -1 \), making the oxidation number of chlorine +7.

The higher the oxidation number, the more oxidized the species is. In our examples, chlorine is more oxidized in perchlorate than in hypochlorite.

Oxidizing Power

The term oxidizing power pertains to the ability of a molecule or ion to oxidize other substances by accepting electrons. Essentially, the higher the oxidation state of an atom within an ion, the greater its oxidizing power. This is because an atom with a high oxidation state tends to more readily accept electrons. For chlorine, evaluating its oxidizing power across different compounds requires looking at its oxidation states:

• In \( \mathrm{ClO}^{-} \), chlorine is in the +1 oxidation state.


• In \( \mathrm{ClO}_{4}^{-} \), the higher oxidation state of +7 suggests a greater oxidizing power compared to \( \mathrm{ClO}^{-} \).

Thus, perchlorate is a stronger oxidizing agent than hypochlorite due to chlorine's higher oxidation state in perchlorate, indicating a higher propensity to accept electrons.

Resonance Structures

Resonance structures provide a way to represent the delocalization of electrons within a molecule which is not adequately described by a single Lewis structure. They are used to demonstrate the multiple ways electrons can be arranged in certain molecules or ions.For ions like the hypochlorite (\( \mathrm{ClO}^{-} \)) and perchlorate (\( \mathrm{ClO}_{4}^{-} \)), resonance structures are important to visualize how chlorine can achieve an octet configuration through the sharing of electrons.

• In \( \mathrm{ClO}^{-} \), the presence of a single bond between Cl and O is typically sufficient.


• In \( \mathrm{ClO}_{4}^{-} \), the formal charge distribution is maintained across four resonance structures with each oxygen atom, which helps to spread out its negative charge and stabilize the structure.

By understanding resonance structures, we can appreciate how the electrons are distributed, which can, in turn, influence the formal charges and subsequent chemical behaviors, although these are distinct from the oxidation states.