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) Perchlorate is a much stronger oxidizing agent than hypochlorite. Suggest an explanation.

Short answer

The formal charge on the chlorine atom in both hypochlorite (\(\mathrm{ClO}^{-}\)) and perchlorate (\(\mathrm{ClO}_{4}^{-}\)) ions is +2. The oxidation numbers of chlorine in \(\mathrm{ClO}^{-}\) and \(\mathrm{ClO}_{4}^{-}\) are +1 and +7, respectively. Perchlorate is a much stronger oxidizing agent than hypochlorite due to the higher oxidation number of chlorine in perchlorate, allowing it to have a greater ability to accept electrons. Additionally, the resonance structures in perchlorate allow for more electron delocalization and ion stability, enhancing its oxidizing ability compared to hypochlorite.

Step-by-step solution

Calculate the formal charge on chlorine in ClO- and ClO4- ions

To determine the formal charge on the chlorine atom in the hypochlorite (ClO-) and the perchlorate (ClO4-) ions, we need to start with the Lewis structures. First, we will calculate the number of valence electrons each ion has. For ClO- ion: Chlorine has 7 valence electrons and oxygen has 6 valence electrons. With the extra electron from the negative charge, the total number of valence electrons is 7 + 6 + 1 = 14. For ClO4- ion: Perchlorate has one chlorine and four oxygen atoms. Thus, the total number of valence electrons will be 7 (for chlorine) + 4 x 6 (for the four oxygens) + 1 (from the negative charge) = 32 electrons. Now, we will find the formal charge on the chlorine atoms in both ions. For ClO- ion: Formal charge on Cl = Valence electrons (in atom) - [Non-bonding electrons + (Bonding electrons / 2)] Formal charge on Cl = 7 - [ 2 + (6 / 2)] = 7 - 5 = +2 For ClO4- ion: There are four resonance structures for ClO4-, in which chorine is double bonded with each oxygen atom one at a time. In each resonance structure, Cl has 2 non-bonding electrons and 6 bonding electrons (from the double bond). Thus the formal charge on Cl in each resonance structure is the same. Formal charge on Cl = 7 - [ 2 + (6 / 2)] = 7 - 5 = +2

Determine the oxidation numbers of chlorine in ClO- and ClO4- ions

To calculate the oxidation number of chlorine in the ClO- and ClO4- ions, we need to assign oxidation numbers to all the atoms in the compounds. For ClO- ion: Since oxygen is more electronegative than chlorine, it will have an oxidation number of -2. For the sum of the oxidation numbers to be equal to the overall charge (-1), the oxidation number of Cl must be +1. For ClO4- ion: As mentioned before, oxygen will have an oxidation number of -2. Therefore, the total negative charge from the four oxygen atoms will be -8. For the sum of the oxidation numbers to be equal to the overall charge (-1), the oxidation number of Cl must be +7.

Explain why perchlorate is a much stronger oxidizing agent than hypochlorite

Perchlorate (ClO4-) is a much stronger oxidizing agent than hypochlorite (ClO-) because the oxidation number of chlorine in perchlorate is higher (+7) than in hypochlorite (+1). A higher oxidation number means that the chlorine in the perchlorate ion has a greater ability to accept electrons, thus making it a stronger oxidizing agent. Additionally, the resonance structures in a perchlorate ion allow more electron delocalization and stability of the ion, further increasing its oxidizing ability compared to hypochlorite.

Key concepts

Formal Charge

Formal charge is a useful concept in chemistry that helps us understand how electrons are distributed in a molecule or ion, and it aids in predicting the structure and reactivity of chemical species. To calculate the formal charge on an atom, we use the formula:

\[\text{Formal charge} = \text{Valence electrons} - (\text{Non-bonding electrons} + \frac{1}{2}\times\text{Bonding electrons})\]

• Valence electrons refer to the electrons present in the outermost shell of the atom in its elemental state.
• Non-bonding electrons are those that are not involved in bond formation.
• Bonding electrons are shared between atoms to form chemical bonds, and each bond is counted twice once for each atom it connects.

For example, in the ClO⁻ ion, chlorine has a formal charge of +2, calculated by knowing it has 7 valence electrons. After accounting for 2 non-bonding and 3 shared (6 from a single bond divided by 2) electrons, we find the formal charge on chlorine to be +2. Observing the formal charge can indicate how atoms maintain their valence shell in a molecule, affecting stability and reactivity.

Resonance Structures

Resonance structures depict various possible configurations of electrons in a molecule, showcasing the concept of electron delocalization. This is crucial in understanding the full picture of how a molecule behaves, as it doesn't always fit into a single, fixed Lewis structure.

When we examine ions like ClO₄⁻, we find multiple resonance structures since the chlorine can form double bonds with the four oxygen atoms in various combinations, each depicting electrons in slightly different locations. Each resonance structure maintains the same arrangement of nuclei and our understanding of electrons can be shared among the atoms.

Key Aspects of Resonance Structures:

• No single structure accurately represents the molecule or ion. Instead, the true structure is a hybrid of all resonance structures.
• Resonance contributes to the stability of a molecule, as electron distribution over multiple atoms allows for lower energy states.
• The resonance structures in perchlorate help explain its stability and effectiveness as a strong oxidizing agent, as electron delocalization reduces energy and increases reactivity potential.

Exploring resonance helps one understand why some molecules are more stable than others, and can predict how molecules will participate in chemical reactions.

Oxidizing Agents

Oxidizing agents play an essential role in chemical reactions. They are substances that accept electrons from other substances, thus undergoing reduction themselves. This electron transfer is fundamental in redox reactions, driving many chemical processes like combustion, corrosion, and metabolism.

In the context of the perchlorate ion (ClO₄⁻), it serves as a potent oxidizing agent because its central chlorine atom has an oxidation number of +7. A higher oxidation number means the atom has a high partially positive character, making it eager to gain electrons and stabilize by achieving a lower oxidation state. Conversely, hypochlorite (ClO⁻) has chlorine at an oxidation number of +1, making it a much less effective oxidizing agent.

• Oxidizing agents are often used in bleaching, sanitizing, and in the synthesis of chemicals.
• They can store energy, such as in batteries, where electron transfer is harnessed to provide power.
• Safety is key, as powerful oxidizing agents can be reactive or even explosive when mixed improperly.

Understanding the strength and function of oxidizing agents like perchlorate informs their many applications, both industrial and lab-scale, driving innovations across chemistry disciplines.