Problem 91 - Chemistry The Central Science In Si Units

Question

The hypochlorite ion, \(\mathrm{ClO}^{-},\) is the active ingredient in bleach. The perchlorate ion, \(\mathrm{ClO}_{4}^{-},\) is a main component of rocket propellants. Draw Lewis structures for both ions. (a) What is the formal charge of Cl in the hypochlorite ion? (b) What is the formal charge of Cl in the perchlorate ion, assuming the Cl-O bonds are all single bonds? (c) What is the oxidation number of Cl in the hypochlorite ion? (d) What is the oxidation number of \(\mathrm{Cl}\) in the perchlorate ion, assuming the Cl- O bonds are all single bonds? (e) In a redox reaction, which ion would you expect to be more easily reduced?

Step-by-Step Solution

Answer

Cl in \(\mathrm{ClO}^-\) has a formal charge of +1, and +3 in \(\mathrm{ClO}_4^-\). Cl oxidation numbers are +1 in \(\mathrm{ClO}^-\) and +7 in \(\mathrm{ClO}_4^-)\). \(\mathrm{ClO}_4^-\) is more easily reduced.

Step 1: Draw Lewis Structure of Hypochlorite Ion

The hypochlorite ion, \(\mathrm{ClO}^-\), consists of one chlorine atom bonded to one oxygen atom with a single bond. Place the negative charge on the more electronegative oxygen atom. This will result in chlorine having 2 lone pairs, while oxygen will have three lone pairs due to the additional negative charge, maintaining the overall charge of the ion as -1. The oxygen atom is more electronegative, and the extra electron will be localized there.

Step 2: Draw Lewis Structure of Perchlorate Ion

The perchlorate ion, \(\mathrm{ClO_4}^-\), comprises a central chlorine atom bonded to four oxygen atoms via single bonds, with a total charge of -1. Consequently, each oxygen atom will have three lone pairs, except one oxygen which will account for the negative charge by possessing an additional lone pair. The chlorine atom usually does not have lone pairs in this structure.

Step 3: Calculate the Formal Charge in Hypochlorite Ion

Use the formula \( \text{Formal Charge} = \text{Valence Electrons} - \text{Non-bonding Electrons} - \frac{1}{2} \times \text{Bonding Electrons} \). For chlorine in \(\mathrm{ClO}^-\), the calculation is: Valence electrons (Cl) = 7, Non-bonding electrons = 4, Bonding electrons = 2.\[ \text{Formal Charge of Cl} = 7 - 4 - \frac{1}{2} \times 2 = +1 \]

Step 4: Calculate the Formal Charge in Perchlorate Ion

Assuming single bonds for \(\mathrm{ClO}_4^-\), calculate: Valence electrons (Cl) = 7, Non-bonding electrons = 0, Bonding electrons = 8 (2 from each Cl-O bond).\[ \text{Formal Charge of Cl} = 7 - 0 - \frac{8}{2} = +3 \]

Step 5: Determine Oxidation Number in Hypochlorite Ion

The oxidation number is calculated based on an assumed complete transfer of electrons:- Oxygen generally has an oxidation number of -2.- For the ion \(\mathrm{ClO}^-\), \(x + (-2) = -1\), solve for \(x\).- This gives Cl an oxidation number of +1.

Step 6: Determine Oxidation Number in Perchlorate Ion

For \(\mathrm{ClO}_4^-\), assuming all \(\mathrm{O}\) atoms have \(-2\), calculate:\(x + 4(-2) = -1\).Solving for \(x\) gives \(x = +7\), thus, the oxidation number of Cl is +7.

Step 7: Determine which Ion is More Easily Reduced

Reduction involves gaining electrons, thus decreasing oxidation number. Since \(\mathrm{ClO}_4^-\) has \(+7\) on Cl (higher oxidation state), it can gain more electrons and is more likely to be reduced than \(\mathrm{ClO}^-\), which has a Cl oxidation number of +1.

Key Concepts

Formal Charge in Lewis Structures

Understanding formal charge is crucial when drawing Lewis structures because it helps us predict the most stable structure for a molecule or ion. Formal charge is not the same as a real charge; it's a bookkeeping tool to assess electron distribution between atoms.
To calculate the formal charge of an atom in a molecule, use the formula:

• Formal Charge = Valence Electrons - Non-bonding Electrons - 0.5 × Bonding Electrons.

In the hypochlorite ion, \(\mathrm{ClO}^-\), chlorine has 7 valence electrons, 4 non-bonding electrons and shares 2 electrons through bonding, giving it a formal charge of +1.
In the perchlorate ion, \(\mathrm{ClO}_{4}^{-}\), chlorine is bonded to four oxygen atoms, sharing 8 electrons. This leads to a formal charge of +3 for chlorine.

Oxidation Number Basics

Oxidation numbers indicate the degree of oxidation of an atom in a compound. These numbers are hypothetical charges assigned based on a set of rules.
Key rules include:

• Oxygen is typically assigned an oxidation number of -2.
• The sum of oxidation numbers in a molecule or ion must equal the total charge.

In the hypochlorite ion, \(\mathrm{ClO}^-\), using the rule that oxygen is -2, chlorine must be +1 to balance the -1 of the ion.
In perchlorate \(\mathrm{ClO}_{4}^{-}\), with each oxygen being -2, chlorine must be +7 to achieve the sum of -1.

Redox Reactions: Gain and Loss of Electrons

Redox reactions involve the transfer of electrons, marked by changes in oxidation numbers. Reduction is the gain of electrons, while oxidation is the loss.
An ion with a higher positive oxidation number is more likely to be reduced because it can accept electrons.
In this context, examine the hypochlorite and perchlorate ions. \(\mathrm{ClO}^-\) has an oxidation number of +1 for chlorine, whereas \(\mathrm{ClO}_{4}^{-}\) has +7.
Since chlorine in \(\mathrm{ClO}_{4}^{-}\) is already at a high oxidation state, it is more likely to gain electrons (be reduced) compared to \(\mathrm{ClO}^-\).

Understanding the Hypochlorite Ion

The hypochlorite ion \(\mathrm{ClO}^-\) is a simple negative ion consisting of a single chlorine atom bonded to an oxygen atom.
It is the active component in common bleach products and acts as an oxidizing agent.
In its Lewis structure, the charge is placed on the more electronegative oxygen, resulting in chlorine having no formal lone pairs.
The oxidation state of chlorine here is +1, indicating it can potentially gain electrons through a reduction reaction.

The Role of Perchlorate Ion in Chemistry

The perchlorate ion \(\mathrm{ClO}_{4}^{-}\) contains a central chlorine atom surrounded by four oxygen atoms. This ion is highly oxidized with the chlorine having an oxidation state of +7.
This high oxidation number indicates it is capable of accepting electrons in a redox reaction, making it a strong oxidizing agent.
Perchlorate is commonly used in applications like rocket propellants due to its ability to release oxygen when decomposed, which aids combustion.
In Lewis structures for \(\mathrm{ClO}_{4}^{-}\), chlorine is found without lone pairs, with each oxygen having three, except one which maintains the negative charge.