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Chapter 6: Problem 18

For Lewis structures, how is the need for multiple bonds generally determined?

Short Answer

Expert verified
Count valence electrons, draw the skeleton, distribute electrons, check the octet rule, and then introduce multiple bonds if needed.

Step by step solution

01

Count the Total Valence Electrons

Sum the valence electrons of all atoms in the molecule. Use the periodic table to determine the number of valence electrons for each atom.
02

Draw the Basic Skeleton Structure

Arrange the atoms with the least electronegative atom in the center and connect each atom using single bonds. Hydrogen is never the central atom.
03

Distribute Remaining Electrons

Distribute the remaining valence electrons as lone pairs starting from the outer atoms, aiming to complete the octet (or duet for hydrogen) for each atom.
04

Check for Octet Rule Violation

If any atom does not satisfy the octet rule after assigning lone pairs, check if the number of electrons matches the total valence electron count.
05

Introduce Multiple Bonds

If there are not enough electrons to satisfy the octet rule for an atom, convert one or more lone pairs from a neighboring atom into a double or triple bond as needed.
06

Verify the Structure

Ensure that all atoms, except hydrogen, follow the octet rule and that the total number of valence electrons is correctly distributed in the structure.

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Key Concepts

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

Valence Electrons
Valence electrons are the electrons located in the outermost shell of an atom. These electrons play a crucial role in chemical bonding. To determine the total valence electron count for a molecule, sum up the valence electrons of each atom involved.
For example, in a water molecule (H₂O), oxygen has 6 valence electrons because it is in Group 16 of the periodic table, and each hydrogen atom has 1 valence electron, resulting in a total of 8 valence electrons for H₂O.
Octet Rule
The octet rule is a principle stating that atoms tend to form bonds in such a way that each atom has eight electrons in its valence shell, achieving a noble gas configuration.
This rule guides the distribution of valence electrons in Lewis structures. However, there are exceptions: hydrogen only needs 2 electrons, while elements like sulfur and phosphorus may have expanded octets when they bond.
Multiple Bonds
Multiple bonds are necessary when there are not enough electrons to satisfy the octet rule for each atom using only single bonds.
Multiple bonds can be double or triple bonds, where two or three pairs of electrons are shared between atoms. For example, in carbon dioxide (CO₂), each oxygen forms a double bond with the central carbon atom to ensure all atoms satisfy their octet.
Electronegativity
Electronegativity is the measure of an atom's ability to attract and hold onto electrons.
When drawing a Lewis structure, the least electronegative atom (excluding hydrogen) is usually placed in the center. Fluorine, oxygen, and nitrogen are some of the most electronegative elements.
Understanding electronegativity helps in determining the molecule's shape and the placement of atoms within a molecule.
Central Atom
The central atom in a Lewis structure is typically the least electronegative atom, except for hydrogen, which is never a central atom.
For instance, in a molecule like NH₃, nitrogen is the central atom because it is less electronegative than hydrogen.
Choosing the correct central atom ensures an accurate Lewis structure, facilitating the correct distribution of remaining valence electrons and satisfying the octet rule where applicable.

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Most popular questions from this chapter

Prepare a report on the work of Linus Pauling. a. Discuss his work on the nature of the chemical bond. b. Linus Pauling was an advocate of the use of vitamin \(\mathrm{C}\) as a preventative for colds. Evaluate Pauling’s claims. Determine if there is any scientific evidence that indicates whether vitamin \(\mathrm{C}\) helps prevent colds.

a. How do ionic and covalent bonding differ? b. How does an ionic compound differ from a molecular compound? c. How does an ionic compound differ from a metal?

Determine the geometry of the following molecules: a. \(\mathrm{CCl}_{4}\) b. \(\mathrm{BeCl}_{2}\) c. \(\mathrm{PH}_{3}\)

Draw a Lewis structure for each of the following polyatomic ions, and then use VSEPR theory to determine the geometry of each: a. \(\mathrm{NO}_{3}^{-}\) b. \(\mathrm{NH}_{4}^{+}\) c. \(\mathrm{SO}_{4}^{2-}\) d. \(\mathrm{ClO}_{2}^{-}\)

a. What is metallic bonding? b. How can the strength of metallic bonding be measured?
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