Book cover for General Chemistry: Principles and Modern Applications

General Chemistry: Principles and Modern Applications

Ralph H. Petrucci, F. Geoffrey Herring, Jeffry D. Madura, Carey Bissonnette

ISBN #9780132931281

11th Edition

3,230 Questions

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293,395 Students Helped

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Summary
Learning Objectives
Key Concepts
Example Problems
Explanations
Common Mistakes

Summary

This chapter integrates multiple models of chemical bonding, starting from Lewis structures and the octet rule to advanced concepts like VSEPR theory, bond order, and electronegativity. It provides a framework for understanding molecular shapes, resonance, and the distinctions between polar and nonpolar bonds, while also introducing foundational ideas in valence bond theory, hybridization, and molecular orbital theory. The chapter emphasizes systematic approaches for writing Lewis structures, predicting molecular geometry, and evaluating bond strengths and energies, which are essential in both theoretical and practical chemistry.

Learning Objectives

1

Describe and draw Lewis structures by applying the octet rule, formal charges, and resonance concepts.

2

Apply VSEPR theory to predict molecular shapes and understand the spatial arrangement of bonds.

3

Analyze the relationship between bond order, bond length, and bond energy in covalent bonding.

4

Differentiate between polar and nonpolar bonds through considerations of electronegativity.

5

Introduce foundational ideas of valence bond theory, hybridization, and molecular orbital theory in chemical bonding.

Key Concepts

CONCEPT

DEFINITION

Lewis Structures

Diagrams that represent the valence electrons of atoms within a molecule, illustrating bonding and lone pairs.

Octet Rule

A chemical rule of thumb that atoms tend to form bonds until they are surrounded by eight electrons in their valence shell.

Formal Charge

A calculated charge on an atom in a molecule, determined by comparing the number of valence electrons in the free atom to those in the bonded state.

Resonance

A situation where more than one valid Lewis structure can be drawn for the same molecule, indicating delocalized electrons.

VSEPR Theory

Valence Shell Electron Pair Repulsion theory, used to predict the three-dimensional molecular shapes based on electron pair repulsions.

Bond Order

The number of chemical bonds between a pair of atoms, which directly affects bond length and bond energy.

Electronegativity

The tendency of an atom to attract electrons (or electron density) towards itself in a chemical bond.

Valence Bond Theory

A theory that describes how atomic orbitals overlap to form chemical bonds in molecules.

Hybridization

The mixing of atomic orbitals to create new hybrid orbitals that can form bonds with appropriate geometries.

Molecular Orbital Theory

A theory that explains the electronic structure of molecules by combining atomic orbitals to form molecular orbitals spread over the entire molecule.

Example Problems

Example 1

Write Lewis symbols for the following atoms. (a) Kr; (b) Ge; (c) N ; (d) Ga; (e) As; (f) Rb.

Example 2

Write Lewis symbols for the following ions. (a) $\mathrm{H}^{-}$; (b) $\operatorname{Sn}^{2+} ;$ (c) $\mathrm{K}^{+} ;$ (d) $\mathrm{Br}^{-} ;$ (e) $\mathrm{Se}^{2-} ;$ (f) $\mathrm{Sc}^{3+}$.

Example 3

Write plausible Lewis structures for the following molecules that contain only single covalent bonds. (a) $\mathrm{FCl} ;$ (b) $\mathrm{I}_{2} ;$ (c) $\mathrm{SF}_{2} ;$ (d) $\mathrm{NF}_{3} ;$ (e) $\mathrm{H}_{2}$ $\mathrm{Te}$.

Example 4

Each of the following molecules contains at least one multiple (double or triple) covalent bond. Give a plausible Lewis structure for (a) $\mathrm{HCN} ;$ (b) $\mathrm{SC}\left(\mathrm{NH}_{2}\right)_{2}$; (c) $\mathrm{F}_{2} \mathrm{CO} ;$ (d) $\mathrm{Cl}_{2} \mathrm{SO} ;$ (e) $\mathrm{C}_{2} \mathrm{H}_{2} ;$ (f) $\mathrm{SO}_{2}$.

Example 5

By means of Lewis structures, represent bonding between the following pairs of elements: (a) $\mathrm{Cs}$ and $\mathrm{Br} ;$ (b) $\mathrm{H}$ and $\mathrm{Sb} ;$ (c) $\mathrm{B}$ and $\mathrm{Cl} ;$ (d) $\mathrm{Cs}$ and $\mathrm{Cl} ;$ (e) $\mathrm{Li}$ and O; (f) $\mathrm{Cl}$ and $\mathrm{I}$. Your structures should show whether the bonding is essentially ionic or covalent.
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Step-by-Step Explanations

QUESTION

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

STEP-BY-STEP ANSWER:

Step 1: Count the total number of valence electrons for all atoms (C has 4, each O has 6; 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 atoms with single bonds and subtract these bonding electrons from the total count (each bond uses 2 electrons).
Step 4: Distribute the remaining electrons to satisfy the octet rule, first assigning electrons to the outer atoms.
Step 5: If outer atoms achieve octet and electrons remain, form double bonds to satisfy the octet rule for the central atom (resulting in O=C=O).
Step 6: Verify that all atoms have completed their octets and that the total number of electrons used equals 16.
Final Answer: The Lewis structure for CO2 is drawn as O=C=O with each oxygen atom having two lone pairs.

Lewis Structure

QUESTION

How do you use VSEPR theory to determine the molecular shape of NH3?

STEP-BY-STEP ANSWER:

Step 1: Draw the Lewis structure for NH3, noting that nitrogen has five valence electrons and each hydrogen contributes one electron.
Step 2: Arrange the atoms with nitrogen in the center and three hydrogen atoms around it, creating three N–H bonds.
Step 3: Identify any lone pairs on the central atom; nitrogen in NH3 has one lone pair.
Step 4: Consider both bonding pairs and lone pairs as electron domains which will repel each other, arranging themselves to minimize repulsion.
Step 5: With four electron domains (three bonds and one lone pair), the electron geometry is tetrahedral.
Step 6: The presence of the lone pair modifies the observed molecular geometry to trigonal pyramidal.
Final Answer: The VSEPR predicted shape for NH3 is trigonal pyramidal.

VSEPR Theory

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Common Mistakes

  • Failing to account for exceptions to the octet rule in molecules that require expanded octets or have electron deficiencies.
  • Incorrectly assigning formal charges by not accurately counting bonding and nonbonding electrons.
  • Misinterpreting resonance structures as separate molecules rather than different representations of the same molecule.
  • Overlooking the role of lone pairs in predicting accurate molecular shapes using VSEPR theory.
  • Confusing polar bonds with polar molecules, without considering the overall molecular geometry and symmetry.