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PowerPoint Lecture Presentation by J
PowerPoint Lecture Presentation by J. David Robertson University of Missouri Chemical Bonding II: Molecular Geometry and Hybridization of Atomic Orbitals Chapter 10 Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
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Valence shell electron pair repulsion (VSEPR) model:
Predict the geometry of the molecule from the electrostatic repulsions between the electron (bonding and nonbonding) pairs. Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry AB2 2 linear linear B 10.1
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HOMEWORK = PREPARE FOR EXAM!!
MODEL SET COLORS BALL COLOR = ATOM TYPE BLUE & BLACK = NON-METALS YELLOW/Red/GREEN = GROUP 1A RED (IIA) = GROUP 11A metals RED (IIIA) = GROUP 111A metals HOMEWORK = PREPARE FOR EXAM!!
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0 lone pairs on central atom 2 atoms bonded to central atom
Cl Be 0 lone pairs on central atom 2 atoms bonded to central atom 10.1
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Arrangement of electron pairs
VSEPR Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry AB2 2 linear linear trigonal planar trigonal planar AB3 3 10.1
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Arrangement of electron pairs
VSEPR Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry AB2 2 linear linear AB3 3 trigonal planar AB4 4 tetrahedral tetrahedral 10.1
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Arrangement of electron pairs
VSEPR Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry AB2 2 linear linear AB3 3 trigonal planar AB4 4 tetrahedral tetrahedral trigonal bipyramidal trigonal bipyramidal AB5 5 10.1
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Arrangement of electron pairs
VSEPR Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry AB2 2 linear linear AB3 3 trigonal planar AB4 4 tetrahedral tetrahedral AB5 5 trigonal bipyramidal AB6 6 octahedral octahedral 10.1
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bonding-pair vs. bonding
pair repulsion lone-pair vs. lone pair repulsion lone-pair vs. bonding >
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Arrangement of electron pairs
VSEPR Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry trigonal planar trigonal planar AB3 3 trigonal planar AB2E 2 1 bent 10.1
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Arrangement of electron pairs
VSEPR Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry AB4 4 tetrahedral tetrahedral trigonal pyramidal AB3E 3 1 tetrahedral 10.1
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Arrangement of electron pairs
VSEPR Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry AB4 4 tetrahedral tetrahedral AB3E 3 1 tetrahedral trigonal pyramidal AB2E2 2 2 tetrahedral bent H O 10.1
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VSEPR trigonal bipyramidal trigonal bipyramidal AB5 5 trigonal
Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry trigonal bipyramidal trigonal bipyramidal AB5 5 trigonal bipyramidal distorted tetrahedron AB4E 4 1 10.1
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VSEPR trigonal bipyramidal trigonal bipyramidal AB5 5 AB4E 4 1
Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry trigonal bipyramidal trigonal bipyramidal AB5 5 AB4E 4 1 trigonal bipyramidal distorted tetrahedron trigonal bipyramidal AB3E2 3 2 T-shaped Cl F 10.1
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VSEPR trigonal bipyramidal trigonal bipyramidal AB5 5 AB4E 4 1
Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry trigonal bipyramidal trigonal bipyramidal AB5 5 AB4E 4 1 trigonal bipyramidal distorted tetrahedron AB3E2 3 2 trigonal bipyramidal T-shaped trigonal bipyramidal AB2E3 2 3 linear I 10.1
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Arrangement of electron pairs
VSEPR Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry AB6 6 octahedral square pyramidal Br F AB5E 5 1 octahedral 10.1
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Arrangement of electron pairs
VSEPR Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry AB6 6 octahedral AB5E 5 1 octahedral square pyramidal square planar Xe F AB4E2 4 2 octahedral 10.1
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Bond Polarity and Polar Molecules
electron rich region electron poor region H F d+ d-
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Predicting Molecular Geometry
Draw Lewis structure for molecule. Count number of electron pairs (bonding & lone) on the central atom and number of atoms bonded to the central atom. Use VSEPR to predict the geometry of the molecule. What are the molecular geometries of SO2 and CF4? F S O F C F C F F S O
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Bond Polarity and Polar Molecules
Polarity refers to have two oppositely charged ends (aka: dipoles) Bond Polarity refers to a specific BOND that has two oppositely charged ends H – F + - or O – C - +
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Bond Polarity and Polar Molecules
Molecular Polarity refers to an ENTIRE MOLECULE having two oppositely charged ends
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Bond Polarity and Polar Molecules
Polar Molecules NEED Polar Bonds! HOWEVER Non-Polar Molecules may have either Polar or Non-Polar bonds Three types of bonds exist between atoms: Non-Polar Covalent Polar Covalent Ionic
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Bond Polarity and Polar Molecules
Non-Polar Covalent Bond Atoms Share Electrons Equally…. H – H Each atoms pulls equally hard on the bonding electrons Bond is essentially neutral…no + or – ends!!
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Bond Polarity and Polar Molecules
2 & 3. Polar Covalent or IONIC Bond Atoms DO NOTShare Electrons Equally…. H – F + - Fluorine atoms pulls harder on the bonding electrons H F
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Bond Polarity and Polar Molecules
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PREDICTING MOLECULAR POLARITY:
STEPS: Draw Lewis Structure as REALISTICALLY as possible. Calculate the ELECTRONEGATIVITY DIFFERENCE between the atoms making the bonds. Determine the TYPES OF BOND using the following scale: In your Lewis Structure draw in the DIRECTION OF ELECTRON SHIFT within the bonds (IF POLAR OR IONIC). Make a judgment (based off symmetry) if the molecule will have a dipole. Polar covalent Ionic Non-polar covalent
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More HW Answers: Part A β = 105 = 109.5 = 120
β = = = 120 φ = = 107 = 180 Part B Ga Ba P N C He Si Br Cs Al Part D Lose electrons Valence electrons, they’re involved in bonding
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More HW Answers: Part D Sharing of valence electrons…polar: e- pulled to one side, non-polar: e- shared equally. Triple: more shell orbitals need to overlap, so atoms move closer together. Triple: more energy required to hold nuclei close together….well compressed spring. EN ≥ 2.0 Already have filled outer-most shell
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More HW Answers: Part D Polar molecules need opposite charges on opposite ends of itself Because + & - ends act like N & S poles In order to give away a valence electron they need to make a bond Part F: Octet, noble Ionic Cation, Anion
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More HW Answers: Part F Sharing NON-POLAR IONIC PULL or HOLD IONS
UNEVENLY ZERO
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More HW Answers: Part G Al2S3 6. NH4Cl Al2(SO4)3 7. MgSO4·5H2O
Al2(SO3) Fe2O3 MgO Na3PO4 Ca(CN) KOH More HW Answers: Part H: S2O P3O4 CO H2O N2O P2O5
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Which of the following molecules have a dipole moment?
H2O, CO2, SO2, and CH4 O H S O dipole moment polar molecule dipole moment polar molecule C H C O no dipole moment nonpolar molecule no dipole moment nonpolar molecule 10.2
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EXAMPLE: HCl STEPS: Draw Lewis Structure as REALISTIC as possible.
Calculate the ELECTRONEGATIVITY DIFFERENCE between the atoms making the bonds. Determine the TYPES OF BOND using the following scale: In your Lewis Structure draw in (if necessary) the DIRECTION OF ELECTRON SHIFT within the bonds. Make a judgment (based off symmetry) if the molecule will have a dipole.
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Does CH2Cl2 have a dipole moment?
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Chemistry In Action: Microwave Ovens
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Sharing of two electrons between the two atoms.
How does Lewis theory explain the bonds in H2 and F2? Sharing of two electrons between the two atoms. Bond Dissociation Energy Bond Length H2 F2 436.4 kJ/mole 150.6 kJ/mole 74 pm 142 pm Overlap Of 2 1s 2 2p Valence bond theory – bonds are formed by sharing of e- from overlapping atomic orbitals. 10.3
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Change in electron density as two hydrogen atoms approach each other.
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Valence Bond Theory and NH3
N – 1s22s22p3 3 H – 1s1 If the bonds form from overlap of 3 2p orbitals on nitrogen with the 1s orbital on each hydrogen atom, what would the molecular geometry of NH3 be? If use the 3 2p orbitals predict 900 Actual H-N-H bond angle is 107.30 10.4
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Hybridization – mixing of two or more atomic orbitals to form a new set of hybrid orbitals.
Mix at least 2 nonequivalent atomic orbitals (e.g. s and p). Hybrid orbitals have very different shape from original atomic orbitals. Number of hybrid orbitals is equal to number of pure atomic orbitals used in the hybridization process. Covalent bonds are formed by: Overlap of hybrid orbitals with atomic orbitals Overlap of hybrid orbitals with other hybrid orbitals 10.4
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Predict correct bond angle 10.4
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Formation of sp Hybrid Orbitals
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Formation of sp2 Hybrid Orbitals
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How do I predict the hybridization of the central atom?
Count the number of lone pairs AND the number of atoms bonded to the central atom # of Lone Pairs + # of Bonded Atoms Hybridization Examples 2 sp BeCl2 3 sp2 BF3 4 sp3 CH4, NH3, H2O 5 sp3d PCl5 6 sp3d2 SF6 10.4
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Sigma bond (s) – electron density between the 2 atoms
Pi bond (p) – electron density above and below plane of nuclei of the bonding atoms Sigma bond (s) – electron density between the 2 atoms 10.5
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Sigma (s) and Pi Bonds (p)
1 sigma bond Single bond Double bond 1 sigma bond and 1 pi bond Triple bond 1 sigma bond and 2 pi bonds How many s and p bonds are in the acetic acid (vinegar) molecule CH3COOH? C H O s bonds = 6 + 1 = 7 p bonds = 1 10.5
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Experiments show O2 is paramagnetic
No unpaired e- Should be diamagnetic Molecular orbital theory – bonds are formed from interaction of atomic orbitals to form molecular orbitals. 10.6
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Energy levels of bonding and antibonding molecular orbitals in hydrogen (H2).
A bonding molecular orbital has lower energy and greater stability than the atomic orbitals from which it was formed. An antibonding molecular orbital has higher energy and lower stability than the atomic orbitals from which it was formed. 10.6
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Molecular Orbital (MO) Configurations
The number of molecular orbitals (MOs) formed is always equal to the number of atomic orbitals combined. The more stable the bonding MO, the less stable the corresponding antibonding MO. The filling of MOs proceeds from low to high energies. Each MO can accommodate up to two electrons. Use Hund’s rule when adding electrons to MOs of the same energy. The number of electrons in the MOs is equal to the sum of all the electrons on the bonding atoms. 10.7
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( ) - bond order = 1 2 Number of electrons in bonding MOs
Number of electrons in antibonding MOs ( - ) bond order 1 10.7
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Delocalized molecular orbitals are not confined between two adjacent bonding atoms, but actually extend over three or more atoms. 10.8
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Electron density above and below the plane of the benzene molecule.
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Chemistry In Action: Buckyball Anyone?
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