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Chapter 8: Problem 81

Describe how formal charges are used to choose between possible molecular structures.

Short Answer

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Question: Explain how to use formal charges to choose between multiple possible molecular structures in a compound. Answer: To use formal charges to choose between multiple possible molecular structures in a compound, follow these steps: 1. Understand the concept of formal charges, which are the apparent charges on individual atoms in a molecule, assuming complete electron transfer in ionic bonds and equal electron sharing in covalent bonds. 2. Calculate the formal charges for each atom in each possible molecular structure using the formula: Formal Charge = (Valence Electrons) - (Non-bonding Electrons + 1/2 Bonding Electrons) 3. Analyze the overall formal charges of the different possible structures. 4. Determine the most stable structure(s) based on the lowest overall formal charge and the presence of negative formal charges on the more electronegative elements. 5. Choose the most stable molecular structure(s) based on the analysis of formal charges and electronegativity.

Step by step solution

01

Understand the concept of formal charges

Formal charges are calculated for each atom within a molecule to help us determine the apparent charge on that atom, based on the assumption that electrons in ionic bonds are completely transferred and that electrons in covalent bonds are equally shared.
02

Calculate the formal charges for each atom in a molecule

To calculate the formal charge of an atom in a molecule, follow this formula: Formal Charge = (Valence Electrons) - (Non-bonding Electrons + 1/2 Bonding Electrons) Valence electrons are the number of electrons in the outer shell of the atom, non-bonding electrons are those that are not involved in any bond, and bonding electrons are those that are shared between atoms in a bond.
03

Analyze multiple possible molecular structures

When given multiple possible molecular structures for a compound, calculate the formal charges for each atom in each structure.
04

Determine the most stable structure(s)

The most stable structures will be those with the lowest overall formal charge, as they are the ones with the electrons distributed most evenly among the atoms. In general, a structure with formal charges closer to zero and with negative formal charges on the more electronegative elements will be more stable.
05

Use formal charges to choose between possible molecular structures

Compare the overall formal charges of the various structures, taking into account the electronegativity of the elements involved, to choose the most stable molecular structure(s).

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Calculating Formal Charges
Understanding the formal charge of an atom within a molecule is essential in predicting the molecular geometry and stability. To calculate the formal charge on an atom, use the formula:

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

Valence electrons are the electrons in the atom's outermost shell, while non-bonding electrons are those not participating in bonds, often found in lone pairs. Bonding electrons are divided by two since they are shared with another atom. The formal charge is a hypothetical charge assuming equal sharing in covalent bonds, an important concept when distinguishing between multiple possible molecular structures. The goal is to have formal charges as close to zero as possible across the molecule, indicating an even distribution of electron density.
Valence Electrons
Valence electrons play a pivotal role in chemical bonding and molecular structure. They are the outermost electrons of an atom and are involved in forming chemical bonds. The number of valence electrons an atom has corresponds to its group number on the periodic table for Group 1 and 2 elements, and 13 through 18 for other elements, minus 10.

For example, carbon is in Group 14, so it has four valence electrons. These valence electrons are used to form covalent bonds with other atoms or exist as lone pairs on an atom. When we study formal charge and molecular stability, it's the valence electrons that we're really focusing on, as they can either be shared in bonds or exist as non-bonding pairs.
Molecular Stability
Molecular stability is influenced by the distribution of electrons across a molecule. A stable molecule typically has a full valence shell for all the atoms involved and minimal strain from bond angles or repulsion between electrons. When calculating formal charges, structures that result in formal charges closest to zero are generally more stable.

Molecular Symmetry and Stability

A symmetrical distribution of electrons leads to a more stable structure as it minimizes the repulsion between negatively charged electrons. Another crucial factor is assigning negative formal charges to more electronegative atoms and positive formal charges to less electronegative ones, which is more natural and leads to a lower energy and more stable configuration.
Electronegativity
Electronegativity refers to the ability of an atom to attract and hold onto electrons. In the context of formal charges, electronegativity helps determine the most stable molecular structure. Atoms with higher electronegativity are better at attracting electrons and will more likely hold a negative formal charge in a stable molecule.

Role in Stability

When comparing possible molecular structures, those with negative formal charges on the more electronegative atoms and positive charges on the less electronegative atoms are generally preferred. The concept of electronegativity allows chemists to predict which atoms within a molecule will remain neutral, which will gain a partial negative charge, and which may end up with a partial positive charge based on their ability to attract electrons.

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Most popular questions from this chapter

A compound with the formula \(\mathrm{Cl}_{2} \mathrm{O}_{6}\) decomposes to a mixture of \(\mathrm{ClO}_{2}\) and \(\mathrm{ClO}_{4} .\) Draw two Lewis structures for \(\mathrm{Cl}_{2} \mathrm{O}_{6}:\) one with a chlorine-chlorine bond and one with a Cl- \(\mathrm{O}-\) Cl arrangement of atoms. Draw a Lewis structure for \(\mathrm{ClO}_{2}\) $$\mathrm{Cl}_{2} \mathrm{O}_{6} \rightarrow \mathrm{ClO}_{2}+\mathrm{ClO}_{4}$$

What is meant by the term polar covalent bond?

If the energy needed to break 2 moles of \(\mathrm{C}=\mathrm{O}\) bonds is greater than the sum of the energies needed to break the \(\mathrm{O}=\mathrm{O}\) bonds in 1 mole of \(\mathrm{O}_{2}\) and vaporize 1 mole of carbon, why does the combustion of pure carbon release heat?

How many electrons are there in the covalent bonds surrounding the central atom in the following species? (a) \(\left(\mathrm{CH}_{3}\right)_{3} \mathrm{Al} ;\) (b) \(\mathrm{B}_{2} \mathrm{Cl}_{4} ;\) (c) \(\mathrm{SO}_{3} ;\) (d) \(\mathrm{SF}_{5}^{-}\).

Oxygen and sulfur combine to form a variety of sulfur oxides. Some are stable molecules and some, including \(\mathrm{S}_{2} \mathrm{O}_{2}\) and \(\mathrm{S}_{2} \mathrm{O}_{3},\) decompose when they are heated. Draw Lewis structures for these two compounds, showing all resonance forms.
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