Book cover for Chemistry The Science in Context

Chemistry The Science in Context

Thomas R. Gilbert

ISBN #9780393615142

5th Edition

2,675 Questions

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191,124 Students Helped

Homework Questions

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Summary
Learning Objectives
Key Concepts
Example Problems
Explanations
Common Mistakes

Summary

Chapter 8 provides a detailed exploration of chemical bonds, emphasizing the importance of understanding the different types of bonds — ionic, covalent, and metallic — and their associated concepts such as Lewis structures, resonance, formal charges, and the relationships between bond order, bond length, and bond energy. The chapter also highlights exceptions to the octet rule, which are critical for predicting molecular behavior and reaction energies. Mastery of these topics not only aids in understanding molecular properties but also in explaining phenomena such as the greenhouse effect in gases.

Learning Objectives

1

Identify and differentiate between the various types of chemical bonds including ionic, covalent (polar and nonpolar), and metallic bonds.

2

Develop skills in drawing accurate Lewis structures, including handling resonance and calculating formal charges to determine the most stable configurations.

3

Understand and apply the relationships between bond order, bond length, and bond energy in predicting molecular properties and reaction enthalpies.

4

Analyze exceptions to the octet rule, including molecules with expanded valence shells and odd-electron configurations.

5

Evaluate how molecular bonding characteristics influence the properties of greenhouse gases.

Key Concepts

CONCEPT

DEFINITION

Chemical Bond

An interaction that holds atoms together in molecules, including ionic, covalent, and metallic bonds.

Ionic Bond

A bond formed through the electrostatic attraction between oppositely charged ions.

Covalent Bond

A bond where atoms share pairs of electrons, which can be further classified as polar or nonpolar depending on electron distribution.

Polar Covalent Bond

A type of covalent bond in which electrons are shared unequally between atoms, leading to partial charges on the atoms.

Metallic Bond

A bond found in metals where electrons are delocalized over a lattice of atoms, contributing to conductivity and malleability.

Lewis Structure

A diagram that represents the bonding between atoms of a molecule and the lone pairs of electrons that may exist.

Resonance

The concept that some molecules cannot be represented by a single Lewis structure but by a hybrid of multiple structures.

Formal Charge

A charge assigned to an atom in a molecule, calculated based on the difference between the number of valence electrons in the free atom and the electrons assigned in the molecule.

Bond Order

The number of chemical bonds between a pair of atoms, related to bond length and bond energy.

Bond Length

The average distance between the nuclei of two bonded atoms; often shorter for bonds with higher bond orders.

Bond Energy

The amount of energy required to break a bond between two atoms, which increases with higher bond order.

Octet Rule

A chemical rule of thumb that atoms tend to form bonds in a way that each atom has eight electrons in its valence shell, resembling a noble gas configuration.

Exceptions to the Octet Rule

Cases where molecules do not or cannot follow the octet rule, such as in molecules with expanded valence shells or odd-electron (radical) configurations.

Example Problems

Example 1

Which group highlighted in Figure P8.1 contains atoms that have the following? (a) 1 valence electron; (b) 4 valence electrons; (c) 6 valence electrons

Example 2

Which of the groups highlighted in Figure P8.2 contains atoms with the following? (a) 2 valence electrons; (b) 3 valence electrons; (c) 5 valence electrons

Example 3

Which of the Lewis symbols in Figure P8.3 correctly portrays the most stable ion of magnesium?

Example 4

What changes must be made to the Lewis symbols in Figure P8.4 to make them correct?

Example 5

Which of the highlighted elements in Figure P8.5 has the greatest bonding capacity?
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Step-by-Step Explanations

QUESTION

How do you draw the Lewis structure for a simple molecule such as CO2?

STEP-BY-STEP ANSWER:

Step 1: Count the total number of valence electrons for all atoms in CO2 (Carbon = 4, Oxygen = 6 each, total = 4 + 6*2 = 16 electrons).
Step 2: Place the least electronegative atom (carbon) in the center and arrange the oxygen atoms symmetrically around it.
Step 3: Connect the atoms with single bonds (2 bonds used, 4 electrons allocated) and subtract these electrons from the total count.
Step 4: Distribute the remaining electrons to complete the octets, starting with the outer atoms (oxygen atoms).
Step 5: If any oxygen does not have an octet, convert lone pairs from oxygen into a double bond with carbon. For CO2, two double bonds are formed.
Step 6: Verify that each atom has a full octet and that the total number of electrons used equals 16.
Final Answer: The Lewis structure of CO2 has a carbon atom double-bonded to two oxygen atoms.

Drawing a Lewis Structure

QUESTION

How do you calculate the formal charge on an atom in a given Lewis structure?

STEP-BY-STEP ANSWER:

Step 1: Identify the number of valence electrons the atom normally has in its free state.
Step 2: Count the number of electrons assigned to the atom in the Lewis structure (each lone pair counts as 2 electrons, and each bonding electron is split equally between the two atoms).
Step 3: Apply the formula: Formal Charge = (Valence electrons) - (Non-bonding electrons) - (Bonding electrons/2).
Step 4: Solve using the values identified. For example, in the nitrate ion (NO3-), calculate the formal charge on one of the oxygen atoms.
Final Answer: The calculated formal charge helps determine which Lewis structure is the most stable.

Calculating Formal Charge

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Common Mistakes

  • Miscounting valence electrons when drawing Lewis structures, leading to incorrect structures.
  • Forgetting to consider exceptions to the octet rule, especially for molecules with expanded valence shells or odd-electron configurations.
  • Confusing resonance structures with a depiction of multiple species, rather than understanding them as contributors to a hybrid structure.
  • Incorrectly assigning formal charges by not properly dividing bonding electrons.
  • Assuming that all chemical bonds behave uniformly, without accounting for the nuanced differences in bond order, bond length, and bond energy.