A Molecular Approach, 4e - Notes for Chapter (9).doc

Chapter 9. Chemical Bonding I: Lewis Theory Chapter 9. Chemical Bonding I: Lewis Theory Chapter 9. Chemical Bonding I: Lewis Theory Student Objectives 9.1 Bonding Models and AIDS Drugs Know that X-ray crystallography was used to characterize the structure of HIV-protease, a biomolecule critical to the reproduction of HIV. Know that Lewis structures are simple predictors of how atoms combine to form ionic compounds and molecules. 9.2 Types of Chemical Bonds Know and understand that chemical bonds form because they lower the potential energy between the charged particles in the constituent atoms. Define and understand ionic bond, covalent bond, and metallic bonding. 9.3 Representing Valence Electrons with Dots Know that valence electrons can be represented with dots around an element symbol. Identify and draw atoms with their valence electrons represented as dots. Know that Lewis theory involves the sharing or transfer of electrons. Define and know the octet rule. 9.4 Ionic Bonding: Lewis Structures and Lattice Energies Draw Lewis structures of ionic compounds containing main-group elements. Understand that the formation of an ionic compound from neutral atoms is exothermic—the amount of energy released is largely due to lattice energy. Know that the Born-Haber cycle is a way of accounting for the energetics of each of the steps in the formation of an ionic compound from its constituent elements. Use a Born-Haber cycle to calculate the lattice energy of an ionic compound. Know that lattice energy decreases for larger ions and increases with increasing charge. Understand why ionic solids are poor electrical conductors while ionic liquids and aqueous solutions of ionic compounds are good electrical conductors. 9.5 Covalent Bonding: Lewis Structures Know that most nonmetal atoms prefer to be surrounded by eight valence electrons, but hydrogen requires only two. Understand that in Lewis theory, a pair of electrons, one from each of two atoms, forms a bond or bonding pair that helps each atom achieve an octet. The two atoms can also share two pairs of electrons (a double bond) or three pairs of electrons (triple bond). Identify and draw covalent compounds with single, double, and triple bonds between constituent atoms. 9.6 Electronegativity and Bond Polarity Know and understand that a pair of electrons does not have to be shared equally between two atoms. Unequal sharing results in a polar covalent bond. Define electronegativity and know its periodic trends. Understand that bonds can range from a nonpolar covalent bond to a polar covalent bond to an ionic bond depending on the difference in electronegativity between the two atoms. Define dipole moment and percent ionic character. 9.7 Lewis Structures of Molecular Compounds and Polyatomic Ions Draw Lewis structures for molecular compounds and polyatomic ions. 9.8 Resonance and Formal Charge Define resonance structures and understand how Lewis structures represent the individual and the hybrid structures. Define formal charge and understand how to calculate it for the atoms in a Lewis structure. 9.9 Exceptions to the Octet Rule: Odd-Electron Species, Incomplete Octets, and Expanded Octets Draw Lewis structures for odd-electron species. Draw Lewis structures for molecules containing atoms with incomplete octets. Draw Lewis structures for molecules containing atoms with expanded octets. Understand why the second-period elements cannot have expanded octets. 9.10 Bond Energies and Bond Lengths Define bond energy. Estimate reaction enthalpies using average bond energies for all bonds broken and formed in a chemical reaction. Understand the inverse relationship between bond length and bond strength. 9.11 Bonding in Metals: The Electron Sea Model Understand how the electron sea model accounts for the general macroscopic properties of metals. Section Summaries Lecture Outline Terms, Concepts, Relationships, Skills Figures, Tables, and Solved Examples Teaching Tips Suggestions and Examples Misconceptions and Pitfalls Lecture Outline Terms, Concepts, Relationships, Skills Figures, Tables, and Solved Examples 9.1 Bonding Models and AIDS Drugs X-ray crystallography structure of HIV-protease Bonding theory Lewis theory Lewis structure Intro figure: HIV-protease and Indinavir unnumbered figure: photo of G. N. Lewis 9.2 Types of Chemical Bonds Types of bonds ionic covalent metallic Electrons in bonds transferred (ionic) shared (covalent) pooled (metallic) Figure 9.1 Ionic, Covalent, and Metallic Bonding unnumbered table: types of bonds Figure 9.2 Possible Configurations of One Negatively Charged Particle and Two Positively Charged Ones 9.3 Representing Valence Electrons with Dots Lewis structure elemental symbol surrounded by valence electrons octet rule Chemical bond sharing electrons octet rule: complete octet by sharing unnumbered figures: Lewis dot structures of period 2 elements Teaching Tips Suggestions and Examples Misconceptions and Pitfalls 9.1 Bonding Models and AIDS Drugs Proteins are complex molecules, but building models of them starts with a single bond. Student understanding of X-ray crystallography can start with a gross simplification that begins with an analogy to X-rays in radiology. In protein structures, X-rays diffract as a result of the nuclei of atoms. Lewis structures were proposed in a J. Am. Chem. Soc. paper in 1916 and have served as a simple but effective bonding model since. 9.2 Types of Chemical Bonds A brief review of electrostatic potential energy should point out that the energy is proportional to the magnitude of the charges and inversely proportional to the distance between them. Figure 9.2 shows the possibilities of one negative charge and two positive charges. Ask the students to imagine possibilities for two negative charges and two positive charges. Conceptual Connection 9.1 Bond Types 9.3 Representing Valence Electrons with Dots The placement of the electrons in the Lewis dot structures for the period 2 elements gives a starting point for students in terms of the number of bonds often formed by these elements. The students will realize with more study that these are not always the case. Lecture Outline Terms, Concepts, Relationships, Skills Figures, Tables, and Solved Examples 9.4 Ionic Bonding: Lewis Structures and Lattice Energies Lewis structures of ionic compounds complete octet by addition or subtraction of electrons Energetics Lattice energy Born-Haber cycle trends ion size ion charge Electrical Conductivity unnumbered figures: Lewis dot structures of KCl and Na2S Example 9.1 Using Lewis Structures to Predict the Chemical Formula of an Ionic Compound Figure 9.3 Lattice Energy Figure 9.4 Born-Haber Cycle for Sodium Chloride unnumbered figure: bond lengths of the group 1A metal chlorides unnumbered figures: bond lengths of NaF and CaO Example 9.2 Predicting Relative Lattice Energies unnumbered figure: melting of NaCl unnumbered figures: electrical conductivity of NaCl(s) and NaCl(aq) Chemistry and Medicine: Ionic Compounds in Medicine 9.5 Covalent Bonding: Lewis Structures Lewis structures of covalent compounds bonding pairs lone pairs single bonds double, triple bonds Covalent bonds (intramolecular) versus intermolecular forces unnumbered figures: Lewis structures for H2O unnumbered figures: Lewis structures for Cl2 and H2 unnumbered figures: Lewis structures for O2, N2 unnumbered figures: Lewis structures for H3O+ and H2O2 Figure 9.5 Intermolecular and Intramolecular Forces Teaching Tips Suggestions and Examples Misconceptions and Pitfalls 9.4 Ionic Bonding: Lewis Structures and Lattice Energies The electron configuration of K+ but not K demonstrates a complete octet. Lattice energy cannot be measured directly in the laboratory but is obtained through a thermodynamic cycle. Ion size and ion charge discussions about the trend in lattice energy are an application of Coulomb’s law. The “reality” section brings focus on the electrical conductivity of ionic compounds. Electricity can be conducted only when the charged particles can move. Conceptual Connection 9.2 Melting Points of Ionic Solids Despite especially the group 1A and group 2A elements’ willingness to ionize, that ionization still is endothermic. Lattice energy is the main reason that ionic solids form from their elements. 9.5 Covalent Bonding: Lewis Structures Introduce the idea of the common numbers of bonds, e.g. 2 for O, 3 for N, 4 for C, etc. It may be useful to use colors or other symbols to differentiate electrons as you assign them in the first few Lewis structure examples. The discussion of intramolecular (microscopic) versus intermolecular (macroscopic) forces is an important one. The latter are always much weaker than bond energies (Section 9.10). Conceptual Connection 9.3 Energy and the Octet Rule Lewis structures and octets are simple schemes to understand how atoms combine to form molecules. They are neither the forces nor the justification for why they do. The overall lower energy of the molecule drives its formation. Lecture Outline Terms, Concepts, Relationships, Skills Figures, Tables, and Solved Examples 9.6 Electronegativity and Bond Polarity Covalent bonds nonpolar covalent polar covalent Electronegativity Pauling values bond polarity, EN covalent, EN = 0–0.4 polar covalent, EN = 0.4–2.0 ionic, EN = 2.0+ Dipole moment = qr percent ionic character Figure 9.6 Orientation of Gaseous Hydrogen Fluoride in an Electric Field Figure 9.7 Electron Density Plot for the HF Molecule Figure 9.8 Electronegativities of the Elements unnumbered figures: molecular models of Cl2, NaCl, and HCl Table 9.1 The Effect of Electronegativity Difference on Bond Type Figure 9.9 Electronegativity Difference (EN) and Bond Type Table 9.2 Dipole Moments of Several Molecules in the Gas Phase Figure 9.10 Percent Ionic Character versus Electronegativity Difference for Some Compounds Example 9.3 Classifying Compounds as Pure Covalent, Polar Covalent, or Ionic 9.7 Lewis Structures of Molecular Compounds and Polyatomic Ions Skeletal structure Total number of valence electrons Distribution of electrons to satisfy octet rule Multiple bonds Example 9.4 Writing Lewis Structures: Write a Lewis Structure for CO2 Example 9.5 Writing Lewis Structures: Write a Lewis Structure for NH3 Example 9.6 Writing Lewis Structures for Polyatomic Ions 9.8 Resonance and Formal Charge Resonance equivalent Lewis structures hybrid structure Formal charge calculation minimization negative formal charges on electronegative atoms unnumbered figure: Lewis structure of O3 Figure 9.11 Hybridization Example 9.7 Writing Resonance Structures unnumbered figures: formal charges of HCN and HNC Example 9.8 Assigning Formal Charges Example 9.9 Drawing Resonance Structures and Assigning Formal Charge for Organic Compounds Teaching Tips Suggestions and Examples Misconceptions and Pitfalls 9.6 Electronegativity and Bond Polarity Polar bonds in diatomic molecules can be detected by placing the molecules in an electric field. Electronegativity has several predictive uses, but differences in electronegativity and the resulting bonds must not be too rigid, particularly values between two types, e.g. 1.9 and 2.0. Conceptual Connection 9.4 Periodic Trends in Electronegativity A dipole moment is a physical measurement, so comparison with the predictions from EN is instructive. Conceptual Connection 9.5 Percent Ionic Character Bonds can be described as a continuum ranging from equal sharing of a pair of electrons between two atoms through degrees of polar covalent bonds to an ionic bond in which the electrons are completely transferred. 9.7 Lewis Structures of Molecular Compounds and Polyatomic Ions Students learn best how to draw Lewis structures from seeing and also working through lots of examples. Common errors in writing good Lewis structures include omitting lone pairs and omitting added or removed electrons in polyatomic ions. 9.8 Resonance and Formal Charge Resonance forms indicate that the actual structure can be represented by several equivalent Lewis structures. The analogy of resonance forms to dogs in Figure 9.11 is one that overcomes the urge to think of the Lewis structures as rapidly changing forms. Formal charges enable students to discriminate among different Lewis structures for a given molecule to determine which is better. Resonance does not arise from two or more interconverting forms. Two or more forms indicate a special situation in which the hybrid structure has special stability or properties, e.g. the different properties of benzene from other alkenes. Formal charges on constituent atoms in a Lewis structure can be confused with oxidation numbers. Lecture Outline Terms, Concepts, Relationships, Skills Figures, Tables, and Solved Examples 9.9 Exceptions to the Octet Rule: Odd-Electron Species, Incomplete Octets, and Expanded Octets Odd-electron species radicals or free-radicals Incomplete octet especially Be, B Expanded octet 10, 12, or even 14 electrons around the central atom impossible for 2nd period elements unnumbered figures: Lewis structures of NO and NO2 Chemistry and the Environment: Free Radicals and the Atmospheric Vacuum Cleaner unnumbered figures: Lewis structures of BF3, BH3, NH3, and BF3NH3 unnumbered figures: Lewis structures of AsF5, SF6, and H2SO4 Example 9.10 Writing Lewis Structures for Compounds Having Expanded Octets 9.10 Bond Energies and Bond Lengths Bond energies breaking bonds: endothermic forming bonds: exothermic higher for multiple bonds average bond energy estimation of Hrxn Bond lengths center-to-center distance between nuclei shorter for multiple bonds unnumbered table: carbon–hydrogen bond energies Table 9.3 Average Bond Energies Figure 9.12 Estimating Hrxn from Bond Energies Example 9.11 Calculating Hrxn from Bond Energies Table 9.4 Average Bond Lengths unnumbered figures: molecular models, bond lengths of halogens unnumbered table: carbon–carbon bond lengths unnumbered table: nitrogen–halogen bond lengths 9.11 Bonding in Metals: The Electron Sea Model Microscopic properties low ionization energies electrons not attached to particular atoms Macroscopic properties thermal and electrical conductivity malleability ductility Figure 9.13 The Electron Sea Model for Sodium Chemistry in the Environment: The Lewis Structure of Ozone unnumbered figure: photo of ductility of Cu Teaching Tips Suggestions and Examples Misconceptions and Pitfalls 9.9 Exceptions to the Octet Rule: Odd-Electron Species, Incomplete Octets, and Expanded Octets Conceptual Connection 9.6 Odd-Electron Species Many radicals exist, especially if one considers organic examples. Incomplete octets are common for boron compounds, but the formal charge for boron when it makes three bonds is 0. Expanded octets are common for structures with P and S as well as for oxyanions of the halogens. Conceptual Connection 9.7 Expanded Octets Expanded octets apply only to central atoms and never to terminal atoms. 9.10 Bond Energies and Bond Lengths Average bond energies are useful for providing approximate values for Hrxn. A bond energy for a bond being broken is positive, while that for a bond being formed is negative. Conceptual Connection 9.8 Bond Energies and Hrxn When estimating Hrxn, students are sometimes tempted to use “products minus reactants” as they did using enthalpies of formation. 9.11 Bonding in Metals: The Electron Sea Model Metals require a unique bonding model to account for many of their macroscopic properties, e.g. electrical conductivity. Additional Problem for Classifying Bonds as Pure Covalent, Polar Covalent, or Ionic (Example 9.3) Determine whether the bond between each pair of atoms will be pure covalent, polar covalent or ionic: a) S and O b) Al and F c) C and Br d) Mg and Cl Solution Find the electronegativity values in Figure 9.10: S = 2.5 O = 3.5 EN = 1.0 polar covalent Solution Find the electronegativity values in Figure 9.10: Al = 1.5 F = 4.0 EN = 2.5 ionic Solution Find the electronegativity values in Figure 9.10: C = 2.5 Br = 2.8 EN = 0.3 pure covalent Solution Find the electronegativity values in Figure 9.10: Mg = 1.2 Cl = 3.0 EN = 1.8 polar covalent, near ionic Additional Problem for Writing Lewis Structures (Examples 9.4, 9.5) Write a Lewis structure for CS2. Write a Lewis structure for CH3Cl. Correct Skeletal Structure Write the correct skeletal structure for the molecule. C is the same electronegativity as S but we put it in the center since there are two S. C is less electronegative than Cl; H is always terminal. Number of Electrons Calculate the total number of electrons for the Lewis structure by summing the valence electrons of each atom in the molecule. C = 4 S = 6 S = 6 total = 16 C = 4 H = 1 H = 1 H = 1 Cl = 7 total = 14 Distribute Electrons Distribute the electrons among the atoms, giving octets (or duets for H) to as many atoms as possible. Begin with the bonding electrons and proceed to lone pairs on the central atom. First, bonding electrons. Next, lone pairs. First, bonding electrons. Next, lone pairs. Form Double of Triple Bonds If any atom lacks an octet, form double or triple bonds as necessary to give them octets. Move lone pairs from S to C. All H have duets. All other atoms have octets. The structure above is complete. Additional Problem for Writing Lewis Structures for Polyatomic Ions (Example 9.6) Write a Lewis structure for NO2. Correct Skeletal Structure Write the correct skeletal structure for the molecule. N is lower electronegativity than O and is at the center. Number of Electrons Calculate the total number of electrons for the Lewis structure by summing the valence electrons of each atom in the molecule. N = 5 O = 6 O = 6 charge = 1 total = 18 Distribute Electrons Distribute the electrons among the atoms, giving octets (or duets for H) to as many atoms as possible. Begin with the bonding electrons and proceed to lone pairs on the central atom. First, bonding electrons. Next, lone pairs. Form Double or Triple Bonds If any atom lacks an octet, form double or triple bonds as necessary to give them octets. Move lone pairs from O to N. Additional Problem for Calculating Hrxn from Bond Energies (Example 9.10) Hydrogen is considered an optimal fuel since it burns without forming any carbon dioxide. 2 H2(g) + O2(g) 2 H2O(g) Use bond energies to calculate Hrxn. Equation Begin by writing the reaction using the Lewis structures of the molecules involved. HH + HH + O=O HOH + HOH Bonds Broken Determine which bonds are broken in the reaction and sum the bond energies of these. Bond broken, H 2 HH = 2(436 kJ) 1 O=O = 1(498 kJ) = 1370 kJ Bonds Formed Determine which bonds are formed in the reaction and sum the negatives of the bond energies of these. Bond formed, H 4 HO = 4(464 kJ) = 1856 kJ Result Find Hrxn by summing the results of the previous two steps. Hrxn = 1370 1856 = 486 kJ 126 Copyright © 2017 by Education, Inc. 125 Copyright © 2017 by Education, Inc.