Chapter 8. Basic Concepts of Chemical Bonding
Media Resources
Figures and Tables in Transparency Pack: Section:
Figure 8.2 Reaction of Sodium Metal with Chlorine 8.2 Ionic Bonding
Gas to Form the Ionic Compound Sodium
Chloride
Table 8.2 Lattice Energies for Some Ionic 8.2 Ionic Bonding
Compounds
Activities: Section:
Periodic Trends: Lewis Symbols 8.1 Lewis Symbols, and the Octet Rule
Octet Rule 8.1 Lewis Symbols, and the Octet Rule
Animations: Section:
H2 Bond Formation 8.3 Covalent Bonding
Periodic Trends: Electronegativity 8.4 Bond Polarity and Electronegativity
Formal Charges 8.5 Drawing Lewis Structures
Movies: Section:
Formation of Sodium Chloride 8.2 Ionic Bonding
3-D Models: Section:
Chlorine 8.1 Lewis Symbols, and the Octet Rule
Basic Concepts of Chemical Bonding
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Other Resources
Further Readings: Section:
The Chemical Bond as an Atomic Tug-of-War 8.1 Lewis Symbols, and the Octet Rule
Gilbert Newton Lewis and the Amazing Electron 8.2 Ionic Bonding
Dots
The Chemical Bond 8.3 Covalent Bonding
Grade-12 Students’ Misconceptions of Covalent 8.3 Covalent Bonding
Bonding and Structure
The Role of Lewis Structures in Teaching Covalent 8.3 Covalent Bonding
Bonding
Electronegativity from Avogadro to Pauling Part II: 8.4 Bond Polarity and Electronegativity
Late Nineteenth- and Early Twentieth-Century
Developments
Demystifying Introductory Chemistry Part 3: 8.4 Bond Polarity and Electronegativity
Ionization Energies, Electronegativity, Polar
Bonds, and Partial Charges
Electron Densities, Atomic Charges, and Ionic, 8.4 Bond Polarity and Electronegativity
Covalent, and Polar Bonds
Drawing Lewis Structures from Lewis Symbols: 8.5 Drawing Lewis Structures
Chapter 8
104
Lewis Structures, Formal Charge, and Oxidation 8.5 Drawing Lewis Structures
Numbers: A More User-Friendly Approach
Valence, Oxidation Number, and Formal Charge: 8.5 Drawing Lewis Structures
Three Related but Fundamentally Different
Concepts
Lost in Lewis Structures: An Investigation of 8.5 Drawing Lewis Structures
Student Difficulties in Developing
Representational Competence
If It’s Resonance, What Is Resonating? 8.6 Resonance Structures
Basic Concepts of Chemical Bonding
105
Chapter 8. Basic Concepts of Chemical Bonding
Common Student Misconceptions
• Students often think that a triple bond is three times as strong as a single bond. The fact that the
second and third bonds ( bonds) are weaker than the first ( bond) needs to be emphasized.
Teaching Tips
• Students need to be able to count the number of valence electrons in order to get the correct Lewis
structure.
• Students need to be reminded that several correct Lewis structures can be often drawn for a molecule
or an ion, but not all correct Lewis structures are equally good.
Lecture Outline
8.1 Lewis Symbols and the Octet Rule
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• The properties of many materials can be understood in terms of their microscopic properties.
• Microscopic properties of molecules include:
• the connectivity between atoms and
• the 3-D shape of the molecule.
• When atoms or ions are strongly attracted to one another, we say that there is a chemical bond
Chapter 8
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The Octet Rule
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• Atoms tend to gain, lose, or share electrons until they are surrounded by eight valence electrons; this
is known as the octet rule.
• We assume that an atom is stable when surrounded by eight electrons (four electron pairs).
FORWARD REFERENCES
• The octet rule will be brought up again in Chapters 22 and 23 for main group nonmetals and
metals, respectively, and in Chapter 24 for carbon.
8.2 Ionic Bonding
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• Consider the reaction between sodium and chlorine:
Na(s) + ½ Cl2(g) → NaCl(s) ∆H°f = –410.9 kJ/mol
• The reaction is violently exothermic.
• We infer that the NaCl is more stable than its constituent elements.
Energetics of Ionic Bond Formation
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• The heat of formation of NaCl(s) is exothermic:
Na(s) + ½ Cl2(g) → NaCl(s) ∆H°f = –410.9 kJ/mol
• Separation of the NaCl into sodium and chloride ions is endothermic:
NaCl(s) → Na+ (g) + Cl–(g) ∆H° = +788 kJ/mol
• where E is the potential energy of the two interacting charged particles, Q1 and Q2 are the
charges on the particles, d is the distance between their centers, and k is a constant:
k = 8.99 109 J-m/C2.
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• As Q1 and Q2 increase, E increases, and as d increases, E decreases.
Electron Configuration of Ions of the s– and p-Block Elements
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• These are derived from the electron configuration of elements with the required number of electrons
Transition-Metal Ions
• Lattice energies compensate for the loss of up to three electrons.
• We often encounter cations with charges of 1+, 2+, or 3+ in ionic compounds.
• However, transition metals can’t attain a noble gas conformation (>3 electrons beyond a noble gas
core).
• Transition metals tend to lose the valence shell electrons first and then as many d electrons as are
required to reach the desired charge on the ion.
• Thus, electrons are removed from 4s before the 3d, etc.
FORWARD REFERENCES
• The link between lattice energy and solubility of ionic compounds will be made in Chapter 17
(section 17.4).
8.3 Covalent Bonding
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• The majority of chemical substances do not have characteristics of ionic compounds.
• We need a different model for bonding between atoms.
• A chemical bond formed by sharing a pair of electrons is called a covalent bond.
• Both atoms acquire noble-gas electronic configurations.
• This is the “glue” to bind atoms together.
Lewis Structures
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• Formation of covalent bonds can be represented using Lewis symbols.
• The structures are called Lewis structures.
• We usually show each electron pair shared between atoms as a line and show unshared electron
pairs as dots.
• Each pair of shared electrons constitutes one chemical bond.
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“Ion Electron Configurations” Activity from Instructor’s Resource CD/DVD
Chapter 8
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• Example: •H + H• → HH has electrons on a line connecting the two H nuclei (H−H).
Multiple Bonds
• It is possible for more than one pair of electrons to be shared between two atoms (e.g., multiple
bonding):
• One shared pair of electrons is a single bond (e.g., H2);
8.4 Bond Polarity and Electronegativity
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• The electron pairs shared between two different atoms are usually unequally shared.
• Bond polarity describes the sharing of the electrons in a covalent bond.
• Two extremes:
• In a nonpolar covalent bond, the electrons are shared equally.
Electronegativity
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• The ability of an atom in a molecule to attract electrons to itself is its electronegativity.
• The electronegativity of an element is related to its ionization energy and electron affinity.
• Pauling electronegativity scale: from 0.7 (Cs) to 4.0 (F).
• Electronegativity increases across a period and decreases down a group.
Electronegativity and Bond Polarity
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• Electronegativity differences close to zero result in nonpolar covalent bonds.
• The electrons are equally or almost equally shared.
• The greater the difference in electronegativity between two atoms, the more polar the bond (polar
covalent bonds).
• There is no sharp distinction between bonding types.
Basic Concepts of Chemical Bonding
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Dipole Moments
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• Molecules like HF have centers of positive and negative charge that do not coincide.
• These are polar molecules.
• We indicate the polarity of molecules in two ways:
Differentiating Ionic and Covalent Bonding
• Interactions of metals and nonmetals often yield ionic compounds.
• When ionic bonding is dominant, we expect compounds to exhibit properties associated with
ionic substances (high-melting solids, strong electrolyte behavior when dissolved in water, etc.).
FORWARD REFERENCES
• Bond polarities combined with molecular shapes (geometries) will be used to determine
molecular polarity in Chapter 9 and physical properties of substances (Chapters 11, 13, 24).
• Polar covalent bonds between atoms of F, O, N, and H will be implicated in hydrogen
bonding in Chapter 11 (section 11.2).
• The electronegativities of atoms in acids will be linked to acid strength in Chapter 16 (section
16.10).
• The electronegativities of nonmetals will be addressed again in sections 22.1 and 22.4.
8.5 Drawing Lewis Structures
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“Molecular Polarity” Activity from Instructor’s Resource CD/DVD
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“Drawing Lewis Structures from Lewis Symbols: A Direct Electron Pairing Approach” from Further
Readings
Chapter 8
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• Some simple guidelines for drawing Lewis structures:
• Sum the valence electrons from all atoms.
• For an anion, add electrons equal to the negative charge.
• For a cation, subtract electrons equal to the positive charge.
Formal Charge and Alternative Lewis Structures
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• Sometimes it is possible to draw more than one Lewis structure with the octet rule obeyed for all
the atoms.
• To determine which structure is most reasonable, we use formal charge.
• The formal charge of an atom is the charge that an atom (in a molecule) would have if all of the
atoms had the same electronegativity.
• For nitrogen:
• There are five valence electrons.
• In the Lewis structure there are two nonbonding electrons and three from the triple bond.
• There are five electrons from the Lewis structure.
• Formal charge = 5 − 5 = 0.
• Using formal charge calculations to distinguish between alternative Lewis structures:
• the most stable structure has the smallest formal charge on each atom and
Basic Concepts of Chemical Bonding
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• the most negative formal charge on the most electronegative atoms.
• It is important to keep in mind that formal charges do NOT represent REAL charges on atoms!
FORWARD REFERENCES
8.6 Resonance Structures
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• Some molecules are not adequately described by a single Lewis structure.
• Typically, structures with multiple bonds can have similar structures with the multiple bonds
between different pairs of atoms.
• Example: Experimentally, ozone has two identical bonds, whereas the Lewis structure
requires one single (longer) and one double bond (shorter).
Resonance in Benzene
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• Benzene belongs to an important category of organic molecules called aromatic compounds.
• Benzene (C6H6) is a cyclic structure.
• It consists of six carbon atoms in a hexagon.
double bonds), we often represent benzene as a hexagon with a circle in it.
FORWARD REFERENCES
• Ozone will be discussed in Chapter 18 (section 18.1).
• Benzene rings will be brought up in Chapter 11 (section 11.7) for liquid crystals and in
Chapter 12 (section 12.8) for graphite and condensation polymers.
• Structure and reactivity of aromatic compounds will be further discussed in Chapter 24.
8.7 Exceptions to the Octet Rule
• There are three classes of exceptions to the octet rule:
• molecules with an odd number of electrons,
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“If It’s Resonance, What Is Resonating?” from Further Readings
Chapter 8
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• molecules in which one atom has less than an octet,
• molecules in which one atom has more than an octet.
Odd Number of Electrons
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• Most molecules have an even number of electrons and complete pairing of electrons occurs.
Less than an Octet of Valence Electrons
• Molecules with less than an octet are also relatively rare.
More than an Octet of Valence Electrons
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• This is the largest class of exceptions.
• Molecules and ions with more than an octet of electrons around the central atom are often called
hypervalent.
electronegative (e.g., F, Cl, O).
FORWARD REFERENCES
• Chemistry of species with an odd number of electrons (radicals) will be discussed in Chapter
18 (section 18.1).
• Electron deficiency of boron will be further discussed in section 22.11.
• Valence shell expansion will be used to explain the formation of noble-gas compounds in
section 22.3.
8.8 Strengths of Covalent Bonds
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• The energy required to break a particular covalent bond in one mole of a gaseous substance is called
the bond enthalpy, D.
• That is, for the Cl2 molecule, D(Cl−Cl) is given by ∆H for the reaction:
Cl2(g) → 2Cl(g).
Basic Concepts of Chemical Bonding
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Bond Enthalpies and the Enthalpies of Reactions
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• We can use bond enthalpies to calculate the enthalpy for a chemical reaction.
• We recognize that in any chemical reaction bonds need to be broken and then new bonds form.
• The enthalpy of the reaction, ∆Hrxn, is given by the sum of bond enthalpies for bonds broken (in
reactants) less the sum of bond enthalpies for bonds formed (in products):
∆Hrxn = ∑D(bonds broken) − ∑D(bonds formed)
Bond Enthalpy and Bond Length
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• The distance between the nuclei of the atoms involved in a bond is called the bond length.
• Multiple bonds are shorter than single bonds.
• We can show that multiple bonds are stronger than single bonds.
• As the number of bonds between atoms increases, the atoms are held closer and more tightly
together.
FORWARD REFERENCES
Chapter 8
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Further Readings:
1. Georgios R. Tsaparlis, “The Chemical Bond as an Atomic Tug–of–War,” J. Chem. Educ., Vol. 61,
1984, 677. An analogy between a covalent bond and a game of tug-and-war is suggested in this
reference.
2001, 1457–1458.
6. Anthony N. Stranges, “Reflections on the Electron Theory of the Chemical Bond: 1900–1925,” J.
Chem. Educ., Vol. 61, 1984, 185–190.
7. William B. Jensen, “Abegg, Lewis, Langmuir, and the Octet Rule,” J. Chem. Educ., Vol. 61, 1984,
191–200.
8. Linus Pauling, “G. N. Lewis and the Chemical Bond,” J. Chem. Educ., Vol. 61, 1984, 201–203.
9. William B. Jensen, “The Use of Dots in Chemical Formulas,” J. Chem. Educ., Vol. 83, 2006, 1590–
1591.
14. Wan-Yaacob Ahmad and Mat B. Zakaria, “Drawing Lewis Structures from Lewis Symbols: A Direct
Electron Pairing Approach,” J. Chem. Educ., Vol. 76, 1999, 329–331.
15. Gordon H. Purser, “Lewis Structures Are Models for Predicting Molecular Structure, Not Electronic
Structure,” J. Chem. Educ., Vol. 76, 1999, 1013–1017.
Basic Concepts of Chemical Bonding
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16. Juan Quilez Pardo, “Teaching a Model for Writing Lewis Structures,” J. Chem. Educ., Vol. 66, 1989,
456–458.
20. Melanie M. Cooper, Nathaniel Gove, Sonia M. Underwood, and Michael W. Klymkowsky, “Lost in
Lewis Structures: An Investigation of Student Difficulties in Developing Representational Competence,”
J. Chem. Educ., Vol. 87, 2010, 869–874.
21. David G. DeWit, “Using Formal Charges in Teaching Descriptive Inorganic Chemistry,” J. Chem.
Educ., Vol. 71, 1994, 750–755.
26. Donald G. Truhlar, “The Concept of Resonance” J. Chem. Educ., Vol. 84, 2007, 781–782.
27. William B. Jensen, “The Origin of the Circle Symbol for Aromaticity” J. Chem. Educ., Vol. 86,
2009, 423-424.
28. R. J. Gillespie, “Electron Densities, Atomic Charges, and Ionic, Covalent, and Polar Bonds,” J.
Chem. Educ., Vol. 78, 2001, 1688–1691.
29. Kenton B. Abel and William M Hemmerlin, “Explaining Resonance–A Colorful Approach,” J.
Chem. Educ., Vol. 68, 1991, 834.
Chapter 8
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32. Solomon H. Snyder and David S. Bredt, “Biological Roles of Nitric Oxide,” Scientific American,
May 1992, 68–77.
Live Demonstrations:
1. Lee. R. Summerlin,, Christie L. Borgford, and Julie B. Ealy, “Bending a Stream of Water,” Chemical
Demonstrations, A Sourcebook for Teachers, Volume 2 (Washington: American Chemical Society, 1988),
p. 91. The polarity of water and cyclohexane are compared in this demonstration.