A Molecular Approach, 4e - Notes for Chapter (10).doc

Chapter 10. Chemical Bonding II Chapter 10. Chemical Bonding II Chapter 10. Chemical Bonding II Student Objectives 10.1 Artificial Sweeteners: Fooled by Molecular Shape Know and understand that energy content and taste are due to microscopic properties related to structure but are independent of each other. Know that taste depends a great deal on the three-dimensional structures of food molecules. Know that a simple model to determine and predict molecular shapes is VSEPR theory, valence shell electron pair repulsion theory. 10.2 VSEPR Theory: The Five Basic Shapes Know and understand that VSEPR theory is based on electron groups that repel each other. Know that VSEPR predicts five basic shapes according to the number of electron groups surrounding a central atom: linear (2), trigonal planar (3), tetrahedral(4), trigonal bipyramidal (5), and octahedral (6). Know the bond angles for each basic shape. Recognize molecules in their correct shapes based on their number of electron groups. 10.3 VSEPR Theory: The Effect of Lone Pairs Understand the difference between electron geometry and molecular geometry. Know and understand the effect of lone pair electrons on molecular geometry with respect to shape and bond angle. Know the different molecular geometries that arise from tetrahedral, trigonal bipyramidal, and octahedral electron geometries. 10.4 VSEPR Theory: Predicting Molecular Geometries Know the procedures for predicting and drawing molecular geometries. Predict and draw the electron and molecular geometries for molecules, including molecules with more than one central atom. 10.5 Molecular Shape and Polarity Identify polar bonds in molecules based on EN. Understand how polar bonds translate into net dipole moments for molecules. Know and understand how vector addition is used to predict net dipole moments. Understand how microscopic polarity results in macroscopic properties of molecules, e.g. the immiscibility of water and oil. 10.6 Valence Bond Theory: Orbital Overlap as a Covalent Bond Understand an interaction energy diagram for the formation of bonds with respect to internuclear distance. Know and understand how the overlap of atomic orbitals leads to bonds and how this is explained by valence bond theory. 10.7 Valence Bond Theory: Hybridization of Atomic Orbitals Define and understand hybridization and the role of atomic orbitals. Know and understand the common types of hybridization: sp3, sp2, and sp. Know the hybridizations for expanded octets: sp3d and sp3d2. Know how to predict hybridization and draw valence bond models of molecules. 10.8 Molecular Orbital Theory: Electron Delocalization Know the basis for molecular orbital theory. Know and understand how linear combinations of atomic orbitals (LCAO) form molecular orbitals. Define bonding orbital and antibonding orbital and understand the differences between the two. Predict and draw molecular orbital diagrams. Understand that molecular orbital theory provides the best explanation of the paramagnetism of O2 and provides the best model for electron delocalization in molecules. 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 10.1 Artificial Sweeteners: Fooled by Molecular Shape Microscopic properties energy from combustion molecular shape sense of taste: active site valence shell electron pair repulsion theory Intro figure: shapes of sugar and aspartame 10.2 VSEPR Theory: The Five Basic Shapes Kinds of electron groups Basic shapes (geometry) linear two electron groups examples: BeCl2, CO2 trigonal planar three electron groups examples: BF3, H2CO tetrahedral four electron groups example: CH4 trigonal bipyramidal five electron groups example: PCl5 octahedral six electron groups example: SF6 Figure 10.1 Repulsion between Electron Groups unnumbered figures: Lewis structures and geometries of BeCl2 and CO2 unnumbered figures: Lewis structures and geometries of BF3 and H2CO Figure 10.2 Representing Electron Geometry with Balloons unnumbered figure: tetrahedral shape unnumbered figure: Lewis structure and geometry of CH4 unnumbered figure: trigonal bipyramidal shape unnumbered figure: Lewis structure and geometry of PCl5 unnumbered figure: octahedral shape unnumbered figure: Lewis structure and geometry of SF6 Example 10.1 VSEPR Theory and the Basic Shapes Teaching Tips Suggestions and Examples Misconceptions and Pitfalls 10.1 Artificial Sweeteners: Fooled by Molecular Shape Molecules have a variety of microscopic and macroscopic properties. Two microscopic properties include energy (heat) released during combustion and shape as interpreted by taste receptors. Ask the students for other potential microscopic properties, e.g. molar mass, solubility in water. What about macroscopic properties, e.g. melting point, density, conductivity? 10.2 VSEPR Theory: The Five Basic Shapes A fundamental idea is that electron groups repel each other. Models are extremely useful for this topic. One can demonstrate the electron groups in class using balloons. Conceptual Connection 10.1 Electron Groups and Molecular Geometry Conceptual Connection 10.2 Molecular Geometry Double bonds can shrink bond angles but otherwise have no effect on geometry. Lecture Outline Terms, Concepts, Relationships, Skills Figures, Tables, and Solved Examples 10.3 VSEPR Theory: The Effect of Lone Pairs Difference between electron and molecular geometries Four electron groups with lone pairs trigonal pyramidal one lone pair smaller bond angle than tetrahedral example: NH3 bent two lone pairs smaller bond angle than tetrahedral or trigonal bipyramidal example: H2O Five electron groups with lone pairs seesaw one lone pair; goes in trigonal plane example: SF4+ T-shaped two lone pairs, both in trigonal plane example: BrF3 linear three lone pairs, all in trigonal plane example: XeF2 Six electron groups with lone pairs square pyramidal example: BrF5 square planar lone pairs 180o apart example: XeF4 unnumbered figure: Lewis structure and geometry of NH3 unnumbered figure: ideal and actual bond angles for NH3 Figure 10.3 Nonbonding versus Bonding Electron Pairs unnumbered figure: Lewis structure and geometry of H2O unnumbered figure: ideal and actual bond angles for H2O Figure 10.4 The Effect of Lone Pairs on Molecular Geometry unnumbered figure: Lewis structure and geometry of SF4 unnumbered figure: Lewis structure and geometry of BrF3 unnumbered figure: Lewis structure and geometry of XeF2 unnumbered figure: Lewis structure and geometry of BrF5 unnumbered figure: Lewis structure and geometry of XeF4 Table 10.1 Electron and Molecular Geometries Teaching Tips Suggestions and Examples Misconceptions and Pitfalls 10.3 VSEPR Theory: The Effect of Lone Pairs Models are extremely useful for showing both kinds of geometry. Conceptual Connection 10.3 Lone Pair Electrons and Molecular Geometry Conceptual Connection 10.4 Molecular Geometry and Electron Group Repulsions Students may be tempted to memorize Table 10.1 (using it for the first few examples is reasonable), but show them how the information can be obtained by logic. Lone pair electrons affect molecular shape. VSEPR theory refers to all electron groups and not just to atoms. Drawing good Lewis structures is essential to predicting geometry. Laziness with Lewis structures leads to omitted lone pairs, leading to incorrect shapes. Students are tempted to place lone pairs at the axes of the trigonal bipyramid rather than in the trigonal plane. Lecture Outline Terms, Concepts, Relationships, Skills Figures, Tables, and Solved Examples 10.4 VSEPR Theory: Predicting Molecular Geometries Procedure Draw Lewis structure. Determine total number of electron groups on central atom. Determine number of bonding groups and number of lone pairs on central atom. Use Table 10.1 to identify electron and molecular geometry. Drawing three-dimensional shapes on paper Shapes of molecules with more than one central atom Examples 10.2 and 10.3 Predicting Molecular Geometries The Nature of Science: Representing Molecular Geometries on Paper unnumbered figure: Lewis structure and geometry of glycine Example 10.4 Predicting the Shape of Larger Molecules 10.5 Molecular Shape and Polarity Polar bonds Net dipole moment Adding dipoles: vector addition one dimension two dimensions three dimensions common cases (Table 10.2) Macroscopic consequences of molecular polarity unnumbered figure: model and electron density plot of HCl unnumbered figure: Lewis structure, model, and electron density plot of CO2 unnumbered figure: Lewis structure, model, and electron density plot of H2O The Nature of Science: Vector Addition Table 10.2 Common Cases of Adding Dipole Moments to Determine whether a Molecule Is Polar Example 10.5 Determining whether a Molecule Is Polar Figure 10.5 Interaction of Polar Molecules unnumbered figure: photo of oil and water unnumbered figure: photo of marbles illustrating interaction of polar molecules with nonpolar molecules Chemistry in Your Day: How Soap Works 10.6 Valence Bond Theory: Orbital Overlap as a Covalent Bond Valence Bond theory interaction energy diagram overlap of atomic orbitals shape determined by geometry of overlapping orbitals Figure 10.6 Interaction Energy Diagram for H2 unnumbered figure: orbital diagram of H and S and orbital-overlap illustration of H2S Teaching Tips Suggestions and Examples Misconceptions and Pitfalls 10.4 VSEPR Theory: Predicting Molecular Geometries Show several examples using the procedure; emphasize the importance of beginning with a good Lewis structure. Introduce the conventions for demonstrating directionality on paper. Shapes of larger molecules are approached by solving the shape for each atom that can be at the center of a set of electron groups. Conceptual Connection 10.5 The Shape of Larger Molecules All pairs of electrons are critical to the procedure. 10.5 Molecular Shape and Polarity Molecular modeling and rendering software (e.g. CHIME, RASMOL, JMOL) can show surface charge maps and structures that can be rotated. Net dipole moments require determining the sum of dipole moments along polar bonds. Review the vector addition rules and examples. Polarity or nonpolarity affects macroscopic properties like boiling point (Ch. 11). Molecules with polar bonds can have a zero net dipole moment. Individual dipoles often add to zero in the basic geometries. 10.6 Valence Bond Theory: Orbital Overlap as a Covalent Bond Electrostatic attraction and repulsion account for the curve when energy is plotted versus internuclear distance in an interaction energy diagram. This section also reviews orbital diagrams from Chapter 8. Conceptual Connection 10.6 What Is a Chemical Bond, Part I? Valence bond theory uses atomic orbital overlap to create bonds in molecules. For some the difficulty may simply involve being able to visualize the overlap, i.e. sort out the portions of the figures. Lecture Outline Terms, Concepts, Relationships, Skills Figures, Tables, and Solved Examples 10.7 Valence Bond Theory: Hybridization of Atomic Orbitals Hybridization and hybrid orbitals sp3 4 equivalent orbitals tetrahedral arrangement sp2 3 equivalent orbitals with a p orbital remaining trigonal planar arrangement sigma bonds and pi bonds rotation restricted for pi bonds sp hybridization 2 equivalent orbitals + 2 p orbitals remaining sp3d hybridization 5 equivalent hybrid orbitals sp3d2 hybridization 6 equivalent hybrid orbitals Determining hybridization and drawing valence bond models Lewis structure VSEPR geometry hybridization molecular sketch hybrid orbitals sigma and pi bonds unnumbered figure: orbital diagram of H and C and geometry of CH4 unnumbered figure: hybridization of carbon’s atomic orbitals Figure 10.7 sp3 hybridization unnumbered figure: valence bond model of CH4 unnumbered figure: valence bond model of NH3 unnumbered figure: orbital diagrams showing sp2 hybridization unnumbered figure: orbital diagrams of H, C, and O and hybridization of carbon’s atomic orbitals Figure 10.8 sp2 Hybridization unnumbered figure: valence bond model of H2CO Figure 10.9 Sigma and Pi Bonding unnumbered figure: valence bond model of H2CO unnumbered figure: valence bond models of 1,2-dichloroethane and 1,2-dichloroethene unnumbered figure: Lewis structures and models of cis-1,2-dichloroethene and trans-1,2-dichloroethene Chemistry in Your Day: The Chemistry of Vision unnumbered figure: orbital diagrams showing sp hybridization Figure 10.10 sp Hybridization unnumbered figure: orbital diagrams of H and C and sp hybridization of C in acetylene unnumbered figure: Lewis structure and valence bond model of C2H2 unnumbered figure: Lewis structure and valence bond model of AsF5 Figure 10.11 sp3d Hybridization Figure 10.12 sp3d2 Hybridization unnumbered figure: Lewis structure and valence bond model of SF6 Table 10.3 Hybridization Scheme from Electron Geometry Examples 10.6 and 10.7 Hybridization and Bonding Scheme Example 10.8 Hybridization and Bonding Scheme Teaching Tips Suggestions and Examples Misconceptions and Pitfalls 10.7 Valence Bond Theory: Hybridization of Atomic Orbitals s and p orbitals are not equivalent, but the combination of s and p creates equivalent hybrid orbitals. Sigma bonds allow free rotation around the bond axis, but pi bonds restrict rotation. The box about the chemistry of vision and retinal is an excellent example. Conceptual Connection 10.7 Number of Hybrid Orbitals Part I Conceptual Connection 10.8 Single and Double Bonds Conceptual Connection 10.9 Number of Hybrid Orbitals Part II Table 10.3 summarizes the hybridization scheme from the electron geometry of hybrid orbitals. Hybrid orbitals are a means to create a set of equivalent orbitals for bonding, but they are not meant to exist until the atoms form the bond. The shapes of the hybrid orbitals may not be obvious from the shapes of the pure s and p orbitals. For example, the small portion of a hybrid orbital may be confusing since it does not appear to get involved with the bond. Lecture Outline Terms, Concepts, Relationships, Skills Figures, Tables, and Solved Examples 10.8 Molecular Orbital Theory: Electron Delocalization Linear combinations of atomic s orbitals constructive: bonding destructive: antibonding Molecular orbital diagrams H2, He2, He2+ bond order Linear combinations of atomic p orbitals shapes of bonding and antibonding orbitals Period 2 homonuclear diatomics 2s-2p mixing paramagnetism and diamagnetism liquid oxygen Period 2 heteronuclear diatomic molecules Polyatomic molecules electron delocalization in ozone, benzene unnumbered figure: bonding MO for H2 unnumbered figure: antibonding MO for H2 Figure 10.13 Formation of Bonding and Antibonding Orbitals unnumbered figure: MO diagram for H2 unnumbered figure: MO diagram for He2 unnumbered figure: MO diagram for He2+ Example 10.9 Bond Order unnumbered figure: MO diagram for Li2 unnumbered figure: MO diagram for Be2 unnumbered figure: sigma bonding and antibonding interactions of 2p orbitals unnumbered figures: pi bonding and antibonding interactions of 2p orbitals unnumbered figure: MO diagram for B2, C2, N2 unnumbered figure: MO diagram for O2, F2, Ne2 Figure 10.14 The Effects of 2s–2p Mixing Figure 10.15 Molecular Orbital Energy Diagrams for Second–Row p–Block Homonuclear Diatomic Molecules unnumbered figure: photo and illustration of O2 paramagnetism unnumbered figure: valence bond model of O2 Example 10.10 Molecular Orbital Theory unnumbered figure: MO diagram for NO Figure 10.16 Shape of 2s Bonding Orbital in NO unnumbered figure: MO diagram for HF Example 10.11 Molecular Orbital Theory for Heteronuclear Diatomic Molecules and Ions unnumbered figure: Lewis structure and valence bond model of O3 unnumbered figure: molecular orbital model of O3 unnumbered figure: resonance structures and molecular orbital model of benzene Teaching Tips Suggestions and Examples Misconceptions and Pitfalls 10.8 Molecular Orbital Theory: Electron Delocalization Molecular orbital theory gives rise to a bonding pair orbital (lower in energy) and a nonbonding orbital (higher in energy). The stability of a molecule or ion composed of two atoms depends on the number of valence electrons that must be used and the orbital energy levels into which those electrons are populated. A stable species must have an overall lower energy relative to that of the constituent atoms. Both sigma and pi have bonding and antibonding orbitals with corresponding appropriate energy levels. Nonbonding molecular orbitals don't help or hinder since they have the same energy as the corresponding atomic orbital. Conceptual Connection 10.10 Bond Order The 2s–2p mixing argument explains the paramagnetism of molecular oxygen. Heteronuclear diatomic molecules demonstrate the effect of different electronegativities and atomic orbital energies on MO formation. MO theory works for more complex molecules and explains the electron delocalization in ozone, benzene, nitrate, and other molecules. Conceptual Connection 10.11 What Is a Chemical Bond, Part II? MO theory is presented from a pictorial point of view, but the details are very mathematical. Oxygen is paramagnetic, illustrating that while Lewis theory is useful, it does have limitations. Additional Problems for Procedure for Predicting Molecular Geometries (Examples 10.2, 10.3) Predict the geometry and bond angles of PCl5. Predict the geometry and bond angles of SOCl2. Lewis Structure Draw a Lewis structure for the molecule. PCl5 has 40 electrons. SOCl2 has 26 electrons. Electron Groups Determine the total number of electron groups around the central atom. Lone pairs, single bonds, double bonds, triple bonds, and single electrons each count as one group. Central atom P has 5 electron groups (single bonds). Central atom S has 4 electron groups (2 single bonds, 1 double bond, 1 lone pair). Bonding Groups and Lone Pairs Determine the number of bonding groups and number of lone pairs around the central atom. Central atom P has… 5 bonding groups 0 lone pairs Central atom S has… 3 bonding groups 1 lone pair Electron and Molecular Geometries Determine the geometries using the list in Table 10.1. The electron geometry is trigonal bipyramidal and molecular geometry is trigonal bipyramidal. Bond angles: 120° equatorial 90° axial The electron geometry is tetrahedral and molecular geometry is trigonal pyramidal. Bond angles: <109.5° Additional Problems for Hybridization and Bonding Scheme (Examples 10.6, 10.7) Write a hybridization and bonding scheme for COCl2. Write a hybridization and bonding scheme for SO42?. Lewis Structure Draw a Lewis structure for the molecule. COCl2 has 24 electrons. SO42? has 32 electrons. Electron Geometry from VSEPR Use VSEPR theory to predict the electron geometry about the central atom. The central atom C has 3 electron groups and therefore a trigonal pyramidal electron geometry. The central atom S has four electron groups and therefore a tetrahedral electron geometry. Central Atom Hybridization Select the correct hybridization for the interior atom(s) based on electron geometry. The trigonal pyramidal electron geometry corresponds to sp2 hybridization. The tetrahedral electron geometry corresponds to sp3 hybridization. Sketch Sketch the molecule, beginning with the central atom and its orbitals. Label Label all bonds using the s or p notation and identify the type of overlapping orbitals. Additional Problem for Bond Order (Example 10.9) Use molecular orbital theory to predict the bond order in HeH+, the cation of a helium and hydrogen atom. Count Electrons Begin by sorting the information in the problem into Given and Find. H: 1 electron He: 2 electrons + charge = 1 electrons total = 2 electrons Energy Diagram HeH+ has two electrons. Assign the electrons to the molecular orbitals, filling lower energy orbitals first and proceeding to higher orbitals. Calculate Bond Order Calculate the bond order by subtracting the number of electrons in antibonding orbitals from the number in bonding orbitals and dividing by two. Result Is it stable? Since the bond order is positive, HeH+ should be stable. Additional Problem for Molecular Orbital Theory (Example 10.10) Draw an MO energy diagram and determine the bond order for O22. Is it diamagnetic or paramagnetic? Energy Level Diagram Write an energy level diagram for the molecular orbitals in O22. Use the energy ordering for O2. Count Valence Electrons Count the valence electrons for the O22 ion. O atom = 6 electrons O atom = 6 electrons 2 charge = 2 electrons total = 14 electrons Fill Molecular Orbitals The O22 ion has 14 valence electrons. Assign the electrons to molecular orbitals with the lowest energy, following Hund’s rule. Calculate Bond Order Calculate the bond order by subtracting the number of electrons in antibonding orbitals from the number in bonding orbitals and dividing by two. Results The bond order is 1. The O22 ion has no unpaired electrons and is neither diamagnetic nor paramagnetic. The ion results from the removal of two protons H+ from hydrogen peroxide H–O–O–H. 140 Copyright © 2017 by Education, Inc. 139 Copyright © 2017 by Education, Inc.