Question

Draw three resonance structures for the chlorate ion \(\left(\mathrm{ClO}_{3}^{-}\right) .\) Show formal charges.

Short answer

Draw three resonance structures moving double bonds among oxygens, keeping net charge -1.

Step-by-step solution

Identify the Total Number of Valence Electrons

The chlorate ion (ClO₃⁻) consists of chlorine (Cl) and three oxygens (O). Chlorine has 7 valence electrons, oxygen has 6 valence electrons each, and the negative charge adds 1 extra electron. Total valence electrons = 7 + (3 × 6) + 1 = 26 electrons.

Draw the Skeleton Structure

Place chlorine in the center and arrange three oxygens around it. Connect each oxygen to chlorine with single bonds. This accounts for 6 electrons.

Distribute Remaining Electrons

Distribute the remaining 20 electrons to complete the octets around the oxygen atoms first (6 electrons each), accounting for 18 electrons (all oxygens now have octets). The remaining 2 electrons will be placed as a lone pair on the chlorine atom.

Determine Formal Charges for the First Structure

Calculate the formal charges using the formula: Formal Charge = Valence electrons - (Non-bonding electrons + 1/2 Bonding electrons) . For central chlorine: 7 - (2+1/2×6) = 0. For each oxygen: 6 - (6+1/2×2) = -1. Thus Chlorine has a formal charge of 0, and each oxygen has -1, yielding a net charge of -1.

Draw Resonance Structures

To draw other resonance forms, move lone pair electrons on chlorine to form double bonds with different oxygens, one at a time, while adjusting electrons to maintain the total count. 1. Convert a lone pair of chlorine and one single bond with one oxygen to a double bond, then repeat similar steps with the other oxygen atoms . Each oxygen bonding with a double bond will have a formal charge of 0 and a singly bonded oxygen will have -1 . Cl remains 0 throughout.

Formal Charges in Resonance Structures

In the final resonance structures, for each that has a Cl=O double bond, that O has 0 formal charge, while other O's remain at -1 . Adjust each structure such that the sum of formal charges remains -1 . For example, in one structure: Cl=O, O^- , O^- where formal charges are +1 on the Cl with double bond oxygen at 0 , and other oxygens at -1.

Key concepts

Chlorate Ion

The chlorate ion, represented as \( \text{ClO}_3^- \), is an anion comprising one chlorine atom and three oxygen atoms. This ion is known for its resonance structures. Resonance structures are different Lewis structures that represent the same ion or molecule, reflecting the delocalization of electrons. In the case of the chlorate ion, resonance helps to explain the actual distribution of electrons.When considering the chlorate ion, it's crucial to note that it has a negative charge. This implies that there is an extra electron added to the total count of valence electrons, making it possible to draw multiple resonance structures while maintaining the overall charge of the ion. Resonance structures are significant in understanding the stability and reactivity of the chlorate ion in various chemical reactions.

Formal Charges

Formal charges offer insight into the charge distribution within a molecule or ion, aiding in identifying the most stable resonance structure. To calculate the formal charge, you use the formula:- \[ \text{Formal Charge} = \text{Valence electrons} - (\text{Non-bonding electrons} + \frac{1}{2} \text{Bonding electrons}) \]- This formula helps you determine whether electrons are shared equally among atoms in a covalent bond.In the chlorate ion \( \text{ClO}_3^- \), the central chlorine has a formal charge of zero in the initial structure, calculated as \( 7 - (2 + \frac{1}{2}\times6) = 0 \). Each oxygen ends up with a formal charge of \(-1\).Balancing formal charges across different resonance structures allows chemists to predict which form contributes most to the resonance hybrid. A structure where atoms have formal charges closer to zero tends to be more stable.

Valence Electrons

Valence electrons must be counted accurately to draw the correct Lewis structures and depict resonance forms for ions like the chlorate ion. Valence electrons are the outermost electrons of an atom and are available for bonding.For chlorate ion \( \text{ClO}_3^- \), the valence electron calculation is crucial. Chlorine has 7 valence electrons, each oxygen has 6, and because of the negative charge, we must add one extra electron. This results in a total of:- \[ 7 + (3 \times 6) + 1 = 26 \text{ valence electrons} \]- These electrons are distributed across the molecule to satisfy the octet rule for each atom whenever possible. This ensures that each atom has a stable configuration in all resonance structures.

Lewis Structures

Lewis structures are diagrams used to represent the valence electrons in an atom, molecule, or ion. They provide a simple way to visualize the arrangement of electrons around atoms.For the chlorate ion \( \text{ClO}_3^- \), the Lewis structure starts with chlorine in the center connected to oxygen atoms with single bonds. After arranging the skeleton structure, we distribute the remaining valence electrons to satisfy the octet rule. This involves placing electrons as lone pairs on oxygen atoms initially, ensuring that each oxygen atom in the skeleton has 8 electrons, either through bond or lone pairs.Once the basic Lewis structure is drawn, resonance structures can be depicted by moving electrons, such as shifting a lone pair from chlorine to form a double bond with oxygen. These adjustments maintain the octet rule while illustrating different potential configurations of electron pairs.Understanding how to draw and interpret Lewis structures is key to visualizing and predicting the behavior and stability of molecules and ions.