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Chapter 9 Outline
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Breakdown, Discussion & Help With Sample Problems: What You Should Know
9.1 Atomic Properties and Chemical Bonds
There are three types of bonds: ionic, covalent, and metallic. Figure 9.2 (p 327) nicely depicts all three--CHEM 200 students should familiarize themselves with all aspects of bonding as presented in this section.
The final paragraphs within Section 9.1 (including Figure 9.3, p 328) introduce Lewis electron-dot symbols for the main group elements in Periods 2 and 3.
Ionic bonds result from the transfer of electrons from metal to nonmetal to form ions that come together to form a solid ionic compound.
The most widely-known ionic compound is sodium chloride (table salt).
click here for a few words about SAMPLE PROBLEM 9.1 (page 329)
One of the more difficult concepts in all of CHEM 200 is that of determining the lattice energy for a given ionic compound. In most cases, the Born-Haber cycle is used to calculate lattice energies.
Lattice energies are the enthalpy changes that occur when gaseous ions coalesce into solid ionic compounds.
Lattice energies result from electrostatic attractions among the oppositely charged ions...and are the sole reason that ionic solids exist because the exothermic nature of the lattice energy process overcomse the energetically unfavorable electron transfers that are required to form the ions.
Fig 9.6 (p 331) depicts the Born-Haber cycle for lithium flouride--wise CHEM 200 students realize the importance of lattice energy determinations and are able to successfully tackle problems such as 9.20 (p 365).
Another important concept discussed in the context of lattice energies is Coulomb's law, which states that the electrostatic force associated with two charges is directly proportional to the product of their magnitudes and inversely proportional to the square of the distance between them...pages 332 and 333 include a discussion of periodic trends in lattice energies, a discussion based in part on Coulomb's law.
9.3 The Covalent Bonding Model
Covalent bonds are much more common than ionic bonds. Covalent bonds are those in which the two electrons within the bond are more-or-less shared equally between the two atoms at either end of the bond.
Fig 9.11 (p 336) is a nice depiction of how the potential energy of a system consisting of two hydrogen atoms varies as the hydrogen atoms approach each other. When the potential energy is at a minimum...that's your bond length between the two hydrogen atoms in H2 (molecular hydrogen). From Silberberg: "It is the mutual attraction of nuclei and electron pair that constitutes a covalent bond." It therefore makes sense that there is an optimal distance for the two hydrogen atoms...in terms of bonding. If the atoms get closer than that optimal distance...the potential energy of the system increases...no doubt due to repulsions between the two positively charged nuclei.
Similarly, if the atoms move to a distance that is further apart than the optimal distance, the nucleus from one is unable to attract the electron of the other...and the energy here, too, increases.
Terms to know in this section include single bond; double bond; triple bond; bond order; bond energy; and bond length.
click here for a few words about SAMPLE PROBLEM 9.2 (page 340)
TOOLS OF THE CHEMISTRY LABORATORY: The graphic that spans pages 342 and 343 includes an informative discussion of Infrared Spectroscopy.
9.4 Between the Extremes: Electronegativity and Bond Polarity
Electronegativity is defined as the relative ability of a bonded atom to attract the shared electrons. All of Section 9.4 is concerned with evaluating the polar character of bonds that aren't purely ionic...but not purely covalent, either. Figs 9.16, 9.17, and 9.18 contain nice descriptions of the raw data...and how those data are used to categorize bonds as mostly ionic; polar covalent; mostly covalent; and nonpolar covalent.
click here for a few words about SAMPLE PROBLEM 9.3 (page 346)
9.5 Depicting Molecules and Ions with Lewis Structures
Drawing structural formulae using the Lewis system is the main emphasis of this section.
CHEM 200 students that have mastered all previous 8+ chapters take to this section like a duck takes to water--important concepts such as the octet rule, valence electrons, resonance and formal charge all become relevant when drawing Lewis Structures.
click here for a few words about SAMPLE PROBLEM 9.4 (page 350)
click here for a few words about SAMPLE PROBLEM 9.5 (page 351)
click here for a few words about SAMPLE PROBLEM 9.6 (page 352)The last section in 9.5 is concerned with the few exceptions to the octet rule...with examples from sections including electron-deficient molecules, odd-electron molecules...and atoms with expanded valence shells.
9.6 Using Lewis Structures and Bond Energies to Calculate Heats of Reaction
Fig 9.21 (p 359) describes the enthalpy changes that occur when reactantsw are converted into products, in terms of evaluating those changes from the perspective of bond energies. Sample Problem 9.9 is a straightforward example of this concept.
click here for a few words about SAMPLE PROBLEM 9.9 (page 361)
9.7 An Introduction to Metallic Bonding
From Silberberg (p 362): "In contrast to ionic bonding, the metal ions are not held in placde as rigidly as in an ionic solid. In contrast to covalent bonding, no particular pair of metal atoms is bonded through any localized pair of electrons. Rather, [in metallic bonding] the valence electrons are shared among all the atoms in the substance, which is held together by the mutual attraction of the metal cations for the mobile, highly delocalized electrons.
Metallic bonding is sometimes referred to as the electron sea model. The electron sea model is used to explain the abnormally high meling points of the Group 2A elements, compared to the Group 1A elements (Fig 9.24, p 363).
exercises such as numbers 9.14, 9.20, 9.25, 9.41, 9.47, 9.52, 9.54, 9.69, and 9.73 are all straightforward problems of the type that are often found on CHEM 200 exams