Question

Assign formal charges to the atoms in the following resonance forms of \(\mathrm{ClO}_{2}^{-}\) :

Short answer

Chlorine has a +1 charge, double-bonded oxygen has a 0 charge, single-bonded oxygen has a -1 charge.

Step-by-step solution

Draw the Lewis Structure

First, we need to draw the Lewis structure for the molecule \(\mathrm{ClO}_{2}^{-}\). The structure can have the chlorine atom as the central atom with two oxygen atoms attached. Remember, the molecule has an extra electron due to its negative charge.

Calculate Valence Electrons

Count the total number of valence electrons available. Chlorine has 7, each oxygen has 6, and there is an extra electron due to the negative charge, totaling to \(7 + 6 + 6 + 1 = 20\) electrons.

Distribute Electrons

Assign electron pairs to form bonds. Start by placing a pair of electrons (single bond) between Cl and each O, using up 4 electrons and leaving 16 electrons to distribute as lone pairs.

Complete Octets and Assign Lone Pairs

Distribute the remaining electrons to complete the octets around each oxygen first. Each oxygen needs 6 more electrons, totaling 12, leaving 4 electrons. These are added as a lone pair to the Cl atom.

Adjust For Resonance and Calculate Charge

Evaluate possible resonance structures by creating double bonds with one of the oxygen atoms. Calculate formal charges for each atom: - Chlorine: \(7 - (2 \text{ bonds} + 4 \text{ lone electrons}) = +1\) or \(7 - (3 \text{ bonds} + 2 \text{ lone electrons}) = 0\)- Oxygen in double bond: \(6 - (4 \text{ bonds}) = 0\)- Single bonded oxygen: \(6 - (1 \text{ bond} + 6 \text{ lone electrons}) = -1\).Each resonance form will differ in the position of the double bond.

Key concepts

Formal Charge

Formal charge is an essential concept in understanding molecules like \(\mathrm{ClO}_{2}^{-}\). It helps determine the most stable resonance structure by assigning charges to atoms based on their electron configuration in a Lewis structure. Simply put, formal charge allows us to see how electrons are shared in a molecule compared to their isolated state.

To compute the formal charge of an atom, use the formula: \[\text{Formal charge} = \text{Valence electrons} - \text{(Bonds + Lone electrons)} \]

For each atom, compare the number of electrons it has when bonded with how many it "owns" in its unbonded, element form.

• Cl in \(\mathrm{ClO}_{2}^{-}\) can have a formal charge of 0 or +1 depending on the bonds it forms.
• Oxygen may have charges based on its bonding and non-bonded electrons.

These calculations clarify the electron distribution and help select the most plausible resonance structure.

Lewis Structure

The Lewis structure serves as the foundation for understanding molecule geometry and bonding. It maps out the arrangement of atoms, bonds, and lone pairs, providing insights into molecular stability and behavior. The first step in drawing a Lewis structure is tallying the valence electrons available, ensuring no electron is overlooked, especially the extra from a charged species like \(\mathrm{ClO}_{2}^{-}\).

Typically, the less electronegative atom sits at the center, surrounded by more electronegative peripheral atoms. For \(\mathrm{ClO}_{2}^{-}\), chlorine is central, with oxygen atoms completing the construct.

Visualizing electron sharing through bonds, and labeling lone pairs are crucial in completing a valid Lewis structure. Always double-check that each atom reaches a stable octet, except for hydrogen or when otherwise specified.

Valence Electrons

Valence electrons play a pivotal role in chemical bonding, marking those electrons that participate in forming bonds. Calculating the correct number of valence electrons forms the backbone of drawing accurate Lewis structures.

In \(\mathrm{ClO}_{2}^{-}\),

• Chlorine (Cl) contributes 7 electrons.
• Each oxygen (O) adds 6 electrons.
• The negative charge implies an additional electron.

Summing these gives us 20 valence electrons for this molecule, which we can allocate to account for bonds and lone pairs.

Understanding these electron contributions aids significantly in properly distributing electrons across the structure, ensuring stability and compliance with molecule geometry standards.

Bonding and Lone Pairs

When discussing bonding and lone pairs, it's about how valence electrons are distributed in molecules to achieve stability. Bonding pairs are electrons shared between atoms to create covalent bonds, while lone pairs are non-bonding electrons residing on a single atom.

For \(\mathrm{ClO}_{2}^{-}\), considering bonding:

• Each chlorine-to-oxygen bond initially uses one pair of electrons.

From the remaining electrons,

• Complete the octets of more electronegative atoms first, typically the peripheral atoms, which in this case are the oxygens.
• Assign any excess onto the central atom, Cl, as lone pairs.

These allocations can shift, forming different resonance structures, which modify how lone pairs and bonding electrons are spread to maintain molecular stability. Understanding this interplay helps visualize different resonance forms and predict molecule behavior more accurately.