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Chapter 8. Basic Concepts of Chemical Bonding

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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 Figure 8.5 Born-Haber Cycle for Formation of NaCl 8.2 Ionic Bonding Figure 8.7 Electronegativity Values Based on 8.4 Bond Polarity and Electronegativity Pauling’s Thermochemical Data Figure 8.8 Electron Density Distribution 8.4 Bond Polarity and Electronegativity Figure 8.11 Oxidation Number, Formal Charge, 8.5 Drawing Lewis Structures and Electron Density Distribution for the HCl Molecule Table 8.4 Average Bond Enthalpies (kJ/mol) 8.8 Strengths of Covalent Bonds Figure 8.15 Using Bond Enthalpies to Calculate 8.8 Strengths of Covalent Bonds ΔHrxn Table 8.5 Average Bond Lengths for Some Single, 8.8 Strengths of Covalent Bonds Double, and Triple Bonds Activities: Section: Periodic Trends: Lewis Symbols 8.1 Lewis Symbols, and the Octet Rule Octet Rule 8.1 Lewis Symbols, and the Octet Rule Coulomb’s Law 8.2 Ionic Bonding Ion Electron Configurations 8.2 Ionic Bonding Molecular Polarity 8.4 Bond Polarity and Electronegativity Writing Lewis Structures I 8.5 Drawing Lewis Structures Writing Lewis Structures II 8.5 Drawing Lewis Structures Bond Enthalpy 8.8 Strengths of Covalent Bonds 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 Phosphorus 8.1 Lewis Symbols, and the Octet Rule Sodium Chloride 8.2 Ionic Bonding Copyright © 2012 Pearson Education, Inc. Basic Concepts of Chemical Bonding 103 Methane 8.5 Drawing Lewis Structures Carbon dioxide 8.5 Drawing Lewis Structures Hydrogen Chloride 8.5 Drawing Lewis Structures Ozone 8.6 Resonance Structures Benzene 8.6 Resonance Structures Phosphorus Pentachloride 8.7 Exceptions to the Octet Rule Chloromethane 8.8 Strengths of Covalent Bonds 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 Reflections on the Electron Theory of the 8.3 Covalent Bonding Chemical Bond: 1900–1925 Abegg, Lewis, Langmuir, and the Octet Rule 8.3 Covalent Bonding G. N. Lewis and the Chemical Bond 8.3 Covalent Bonding The Use of Dots in Chemical Formulas 8.3 Covalent Bonding Electronegativity and Bond Type: Predicting Bond 8.4 Bond Polarity and Electronegativity Type Electronegativity from Avogadro to Pauling 8.4 Bond Polarity and Electronegativity Part I: Origins of the Electronegativity Concept 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 A Direct Electron Pairing Approach Lewis Structures Are Models for Predicting 8.5 Drawing Lewis Structures Molecular Structure, Not Electronic Structure Teaching a Model for Writing Lewis Structures 8.5 Drawing Lewis Structures Drawing Lewis Structures without Anticipating 8.5 Drawing Lewis Structures Octets The ‘6N + 2 Rule’ for Writing Lewis Octet 8.5 Drawing Lewis Structures Structures Another Procedure for Writing Lewis Structures 8.5 Drawing Lewis Structures Using Formal Charges in Teaching Descriptive 8.5 Drawing Lewis Structures Inorganic Chemistry Copyright © 2012 Pearson Education, Inc. 104 Chapter 8 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 Aromatic Bagels: An Edible Resonance Analogy 8.6 Resonance Structures The Concept of Resonance 8.6 Resonance Structures Explaining Resonance—A Colorful Approach 8.6 Resonance Structures A Visual Aid for Teaching the Resonance Concept 8.6 Resonance Structures The Origin of the Circle Symbol for Aromaticity 8.6 Resonance Structures Nitric Oxide—Some Old and New Perspectives 8.7 Exceptions to the Octet Rule Biological Roles of Nitric Oxide 8.7 Exceptions to the Octet Rule The Relative Explosive Power of Some Explosives 8.8 Strengths of Covalent Bonds Exothermic Bond Breaking: A Persistent 8.8 Strengths of Covalent Bonds Misconception Copyright © 2012 Pearson Education, Inc. 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 (π bonds) are weaker than the first (σ bond) needs to be emphasized. • Students confuse formal charges with real charges on atoms. • The only place the ↔ arrow is used is for resonance; students often want to use this to indicate equilibrium. • Students do not appreciate that the exceptions to the octet rule are almost as common as the examples of substances that obey it. • Students often confuse the octet rule with having any 8 valence electrons (e.g., 4s2 3d6 for iron). • Students often think that all polar substances conduct electricity. 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 Rule1,2,3 • 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 between them. • In chemical bonds, electrons are shared or transferred between atoms. • Types of chemical bonds include: • ionic bonds (electrostatic forces that hold ions together, e.g., NaCl); • covalent bonds (result from sharing electrons between atoms, e.g., Cl2); • metallic bonds (refers to metal nuclei floating in a sea of electrons, e.g., Na). • The electrons involved in bonding are called valence electrons. • Valence electrons are found in the incomplete, outermost shell of an atom. • As a pictorial understanding of where the electrons are in an atom, we represent the electrons as dots around the symbol for the element. • The number of valence electrons available for bonding are indicated by unpaired dots. • These symbols are called Lewis symbols or Lewis electron-dot symbols. • We generally place the electrons on four sides of a square around the element’s symbol. 1 “The Chemical Bond as an Atomic Tug-of-War” from Further Readings 2 “Chlorine” 3-D Model from Instructor’s Resource CD/DVD 3 “Phosphorus” 3-D Model from Instructor’s Resource CD/DVD Copyright © 2012 Pearson Education, Inc. 106 Chapter 8 The Octet Rule4 • Atoms tend to gain, lose, or share electrons until they are surrounded by eight valence electrons; this is known as the octet rule. • An octet consists of full s and p subshells. • We know that ns2np6 is a noble gas configuration. • 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 Bonding5,6,7,8 • Consider the reaction between sodium and chlorine: Na(s) + ½ Cl2(g) Æ 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. • Sodium has lost an electron to become Na+ and chlorine has gained the electron to become Cl–. • Note that Na+ has an Ne electron configuration and Cl– has an Ar configuration. • That is, both Na+ and Cl– have an octet of electrons. • NaCl forms a very regular structure in which each Na+ ion is surrounded by six Cl– ions. • Similarly, each Cl– ion is surrounded by six Na+ ions. • There is a regular arrangement of Na+ and Cl–. • Note that the ions are packed as closely as possible. • Ionic substances are often crystalline, brittle compounds with high melting points. Energetics of Ionic Bond Formation9,10,11 • The heat of formation of NaCl(s) is exothermic: Na(s) + ½ Cl2(g) Æ NaCl(s) ∆H°f = –410.9 kJ/mol • Separation of the NaCl into sodium and chloride ions is endothermic: NaCl(s) Æ Na+ (g) + Cl–(g) ∆H° = +788 kJ/mol • The energy required to separate one mole of a solid ionic compound into gaseous ions is called the lattice energy, ∆Hlattice. • Lattice energy depends on the charge on the ions and the size of the ions. • The stability of the ionic compound comes from the attraction between ions of unlike charge. • The specific relationship is given by Coulomb’s equation: Q1Q2 E = k d • 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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