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.



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Chapter 10. Chemical Bonding II 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 O 2 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



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Chapter 10. Chemical Bonding II Lecture Outline Terms, Concepts, Relationships, Skills 10.1 Artificial Sweeteners: Fooled by Molecular Shape  Microscopic properties o energy from combustion o molecular shape  sense of taste: active site  valence shell electron pair repulsion theory 10.2 VSEPR Theory: The Five Basic Shapes  Kinds of electron groups  Basic shapes (geometry) o linear  two electron groups  examples: BeCl 2 , CO 2 o trigonal planar  three electron groups  examples: BF 3 , H 2 CO o tetrahedral  four electron groups  example: CH 4 o trigonal bipyramidal  five electron groups  example: PCl 5 o octahedral  six electron groups  example: SF 6



Figures, Tables, and Solved Examples

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Intro figure: shapes of sugar and aspartame

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Figure 10.1 Repulsion between Electron Groups unnumbered figures: Lewis structures and geometries of BeCl 2 and CO 2 unnumbered figures: Lewis structures and geometries of BF 3 and H 2 CO Figure 10.2 Representing Electron Geometry with Balloons unnumbered figure: tetrahedral shape unnumbered figure: Lewis structure and geometry of CH 4 unnumbered figure: trigonal bipyramidal shape unnumbered figure: Lewis structure and geometry of PCl 5 unnumbered figure: octahedral shape unnumbered figure: Lewis structure and geometry of SF 6 Example 10.1 VSEPR Theory and the Basic Shapes

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Chapter 10. Chemical Bonding II Teaching Tips Suggestions and Examples 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





Misconceptions and Pitfalls

Double bonds can shrink bond angles but otherwise have no effect on geometry.

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Chapter 10. Chemical Bonding II Lecture Outline Terms, Concepts, Relationships, Skills 10.3 VSEPR Theory: The Effect of Lone Pairs  Difference between electron and molecular geometries  Four electron groups with lone pairs o trigonal pyramidal  one lone pair  smaller bond angle than tetrahedral  example: NH 3 o bent  two lone pairs  smaller bond angle than tetrahedral or trigonal bipyramidal  example: H 2 O  Five electron groups with lone pairs o seesaw  one lone pair; goes in trigonal plane  example: SF 4 + o T‐shaped  two lone pairs, both in trigonal plane  example: BrF 3 o linear  three lone pairs, all in trigonal plane  example: XeF 2  Six electron groups with lone pairs o square pyramidal  example: BrF 5 o square planar  lone pairs 180o apart  example: XeF 4

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Figures, Tables, and Solved Examples unnumbered figure: Lewis structure and geometry of NH 3 unnumbered figure: ideal and actual bond angles for NH 3 Figure 10.3 Nonbonding versus Bonding Electron Pairs unnumbered figure: Lewis structure and geometry of H 2 O unnumbered figure: ideal and actual bond angles for H 2 O Figure 10.4 The Effect of Lone Pairs on Molecular Geometry unnumbered figure: Lewis structure and geometry of SF 4 + unnumbered figure: Lewis structure and geometry of BrF 3 unnumbered figure: Lewis structure and geometry of XeF 2 unnumbered figure: Lewis structure and geometry of BrF 5 unnumbered figure: Lewis structure and geometry of XeF 4 Table 10.1 Electron and Molecular Geometries



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Chapter 10. Chemical Bonding II Teaching Tips Suggestions and Examples 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.

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Misconceptions and Pitfalls 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.

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Chapter 10. Chemical Bonding II Lecture Outline Terms, Concepts, Relationships, Skills 10.4 VSEPR Theory: Predicting Molecular Geometries  Procedure o Draw Lewis structure. o Determine total number of electron groups on central atom. o Determine number of bonding groups and number of lone pairs on central atom. o 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 10.5 Molecular Shape and Polarity  Polar bonds  Net dipole moment  Adding dipoles: vector addition o one dimension o two dimensions o three dimensions o common cases (Table 10.2)  Macroscopic consequences of molecular polarity

Figures, Tables, and Solved Examples

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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

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10.6 Valence Bond Theory: Orbital Overlap as a Covalent Bond  Valence Bond theory o interaction energy diagram o overlap of atomic orbitals o shape determined by geometry of overlapping orbitals



 

unnumbered figure: model and electron density plot of HCl unnumbered figure: Lewis structure, model, and electron density plot of CO 2 unnumbered figure: Lewis structure, model, and electron density plot of H 2 O 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 Figure 10.6 Interaction Energy Diagram for H 2 unnumbered figure: orbital diagram of H and S and orbital‐overlap illustration of H 2 S

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Chapter 10. Chemical Bonding II Teaching Tips Suggestions and Examples 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. 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). 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 10.5 What Is a Chemical Bond, Part I?



Misconceptions and Pitfalls



All pairs of electrons are critical to the procedure.

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Molecules with polar bonds can have a zero net dipole moment. Individual dipoles often add to zero in the basic geometries.

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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.

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Chapter 10. Chemical Bonding II Lecture Outline Terms, Concepts, Relationships, Skills 10.7 Valence Bond Theory: Hybridization of Atomic Orbitals  Hybridization and hybrid orbitals o sp3  4 equivalent orbitals  tetrahedral arrangement o sp2  3 equivalent orbitals with a p orbital remaining  trigonal planar arrangement o sigma bonds and pi bonds  rotation restricted for pi bonds o sp hybridization  2 equivalent orbitals + 2 p orbitals remaining o sp3d hybridization  5 equivalent hybrid orbitals 3 2 o sp d hybridization  6 equivalent hybrid orbitals  Determining hybridization and drawing valence bond models o Lewis structure o VSEPR geometry o hybridization o molecular sketch  hybrid orbitals  sigma and pi bonds



Figures, Tables, and Solved Examples                         

unnumbered figure: orbital diagram of H and C and geometry of CH 4 unnumbered figure: hybridization of carbon’s atomic orbitals Figure 10.7 sp3 hybridization unnumbered figure: valence bond model of CH 4 unnumbered figure: valence bond model of NH 3 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 H 2 CO Figure 10.9 Sigma and Pi Bonding unnumbered figure: valence bond model of H 2 CO 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 Figure 10.10 sp Hybridization unnumbered figure: orbital diagrams showing 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 C 2 H 2 unnumbered figure: Lewis structure and valence bond model of AsF 5 Figure 10.11 sp3d Hybridization Figure 10.12 sp3d2 Hybridization unnumbered figure: Lewis structure and valence bond model of SF 6 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



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Chapter 10. Chemical Bonding II Teaching Tips Suggestions and Examples 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 10.6 Single and Double Bonds  Table 10.3 summarizes the hybridization scheme from the electron geometry of hybrid orbitals.

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Misconceptions and Pitfalls 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.

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Chapter 10. Chemical Bonding II Lecture Outline Terms, Concepts, Relationships, Skills 10.8 Molecular Orbital Theory: Electron Delocalization  Linear combinations of atomic s orbitals o constructive: bonding o destructive: antibonding o Molecular orbital diagrams  H 2 , He 2 , He 2 +  bond order  Linear combinations of atomic p orbitals o shapes of bonding and antibonding orbitals o Period 2 homonuclear diatomics o

o o

2s‐2p mixing  paramagnetism and diamagnetism  liquid oxygen Period 2 heteronuclear diatomic molecules Polyatomic molecules  electron delocalization in ozone, benzene

Figures, Tables, and Solved Examples                         

unnumbered figure: bonding MO for H 2 unnumbered figure: antibonding MO for H 2 Figure 10.13 Formation of Bonding and Antibonding Orbitals unnumbered figure: MO diagram for H 2 unnumbered figure: MO diagram for He 2 unnumbered figure: MO diagram for He 2 + Example 10.9 Bond Order unnumbered figure: MO diagram for Li 2 unnumbered figure: MO diagram for Be 2 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 B 2 , C 2 , N 2 unnumbered figure: MO diagram for O 2 , F 2 , Ne 2 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 O 2 paramagnetism unnumbered figure: valence bond model of O 2 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 O 3 unnumbered figure: molecular orbital model of O 3 unnumbered figure: resonance structures and molecular orbital model of benzene



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Chapter 10. Chemical Bonding II Teaching Tips Suggestions and Examples 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.  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 10.7 What Is a Chemical Bond, Part II?

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Misconceptions and Pitfalls 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.

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Chapter 10. Chemical Bonding II Additional Problems for Procedure for Predicting Molecular Geometries (Examples 10.2, 10.3)

Predict the geometry and bond angles of PCl 5 .

Lewis Structure

PCl 5

has 40 electrons.

Predict the geometry and bond angles of SOCl 2 .

26

SOCl 2 has electrons.

Draw a Lewis structure for the molecule.

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).

Central atom P has…

Central atom S has…

5 bonding groups

3 bonding groups

0 lone pairs

1 lone pair

Bonding Groups and Lone Pairs Determine the number of bonding groups and number of lone pairs around the central atom.

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.

The electron geometry is tetrahedral and molecular geometry is trigonal pyramidal.

o Bond angles: 120 equatorial

Bond