Chemical Bonding II: Molecular Geometry and Hybridization of Atomic Orbitals Chapter 10

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Valence Shell Electron-Pair Repulsion (VSEPR) Theory Molecular Shapes Valence Shell Electron-Pair Repulsion (VSEPR) Theory Each group of valence electrons around a central atom is located as far away from the others as possible in order to minimize repulsion. An electron group may be a single, double, or triple bond or a lone pair.

Balloon Analogy for the Mutual Repulsion of Electron Groups Two Three Four Five Six Number of Electron Groups

First we will look at examples where there are no lone pairs of electrons on the central atom. In this case the molecular geometry is the same as the electron geometry.

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

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

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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Chapter 10 Homework 1-5, 7, 8, 10-14, 16-24, 26, 28, 29, 32, 34-36, 41, 42

Last Test – Dec 6 or 7 EH Assignment Due Nov 29 Unit 12 Polyatomic Structures Minimum score Sec 1 Electronegativity and Bond Polarity 90 Sec 2 Lewis Dot Diagrams 88 Sec 3 Shapes of Ions and Molecules later Sec 4 Resonance and Formal Charge Sec 5 Review of Molecular Shapes Last Test – Dec 6 or 7

Molecular Shapes You must distinguish between the electron-group arrangement and the molecular shape. The molecular shape is defined by the relative positions of the atomic nuclei. If there are no lone pairs of electrons the molecular shape is the same as the electron-group arrangement. If there are lone pairs the molecular shape is only part of the electron arrangement. The ideal bond angles are determined by the electron-group angles.

bonding-pair vs. bonding pair repulsion lone-pair vs. lone pair repulsion lone-pair vs. bonding >

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 SO2 10.1

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

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

Arrangement of electron pairs VSEPR 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 seesaw AB4E 4 1 SF4 10.1

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

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

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

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

Predicting Molecular Geometry Draw Lewis structure for molecule. Count number of lone pairs 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 AsH3, OF2, AlCl4-, H3O+, C2H4, SnCl5-, and C2H2? 10.1

Molecular Shapes with More Than One Central Atom For more complicated molecules you must give the shape and bond angles around each central atom. Examples: ethanol

The Tetrahedral Centers of Ethane and of Ethanol

Polar Molecules - Dipole moments In chapter 9 you learned about covalent bonds that are polar. A polar molecule is one the has a positive end and a negative end. (The center of positive charge and the center of negative charge do not coincide.) A polar molecule is said to have a dipole moment. A molecule with two or more polar bonds can be nonpolar if it is symmetric. Dipole moment = charge x distance

Dipole Moments and Polar Molecules H F electron rich region electron poor region d+ d- m = Q x r Q is the charge r is the distance between charges 1 D = 3.36 x 10-30 C m 10.2

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- - - - .. .. .. .. .. .. Polarity of CO2 and H2O + O O C O H H + + Water - H2O Carbon Dioxide - CO2 - - - + - .. .. .. .. O O C O .. .. H H + + A non- polar molecule A Polar molecule The bonds are polar, but the molecule is symmetrical, so that the molecule overall is non-polar. The bonds are polar, and the molecule is non-symmetrical.

.. .. .. Predicting The Polarity of Molecules Problem: From electronegativity, and their periodic trends, predict whether each of the following molecules is polar and show the direction of bond dipoles and the overall molecular dipole. (a) Phosphine, PH3 (b) Carbon disulfide, CS2 (also OCS) (c) Aluminum chloride, AlCl3 .. .. .. (a) P P P H H H H H H H H H Molecular Dipole Bond dipoles Molecular shape

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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

Which of these are polar? CH4, CCl4, CHCl3,, CH2Cl2, CH3Cl. 10.2

Valence Bond Theory Basic Principle of Valence Bond Theory: a covalent bond forms when the orbitals from two atoms overlap and a pair of electrons occupies the region between the nuclei. Usually this means that each bonding orbital should contain one electron. The electrons must have opposite spins. The bond strength depends on the attraction of nuclei for the shared electrons, so the greater the orbital overlap, the stronger the bond.

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Change in electron density as two hydrogen atoms approach each other. 10.3

The sp Hybrid Orbitals in Gaseous BeCl2

The total number of orbitals doesn’t change. Hybridization is the mixing of atomic orbitals in an atom to generate a set of new orbitals called hybrid orbitals. The total number of orbitals doesn’t change. Covalent bonds are formed by: Overlap of hybrid orbitals with atomic orbitals Overlap of hybrid orbitals with other hybrid orbitals

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How do I predict the hybridization of the central atom? Count the groups of electrons and determine the electron geometry. # electron groups + Electron geometry Hybridization Examples 2 (linear) sp BeCl2 3 (trigonal planar) sp2 BF3 4 (tetrahedral) sp3 CH4, NH3, H2O 5 (trig. bipyramid) sp3d PCl5 6 (octahedral) sp3d2 SF6 10.4

The sp3d Hybrid Orbitals in PCl5

The sp3d2 Hybrid Orbitals in SF6 Sulfur Hexafluoride -- SF6

Predict correct bond angle

All the bonds we have seen so far are called sigma bonds: the highest electron density is along the bond axis. All single bonds are sigma bonds and the first bond in double or triple bonds is a sigma bond. A pi bond is formed by a side-to-side overlap of p orbitals. There are regions of high electron density above and below the bond axis. The second bond of a double bond is a pi bond, as are two of the bonds in a triple bond.

Describe the bonding in C2H4 10.5

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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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For Chapter 10 See Test 1 Spring 2001 Test 5 1998

Last Test – EH Assignment Due Dec 1 Unit 12 Polyatomic Structures Minimum score Sec 1 Electronegativity and Bond Polarity - Sec 2 Lewis Dot Diagrams Sec 3 Shapes of Ions and Molecules 90 Sec 4 Resonance and Formal Charge Sec 5 Review of Molecular Shapes Last Test –

Last Test – EH Assignment Due Dec 2 Unit 13 Covalent Bonding Minimum score Sec 1 Orbital Shapes 90 Sec 2 Hybridization 85 Sec 3 Bonding in Organic Structures Sec 4 Polarity of Molecules Sec 5 Molecular Orbital Structures optional Last Test –

Describe the bonding in C2H2 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

Skip pages 318-326 molecular orbitals Resonance implies electron delocalization. This can be explained by valence bond and molecular orbital theory. This occurs when a p orbital overlaps with more than one other p orbital. The  electrons are free to move around three or more atoms. Examples: ozone, benzene, sulfur trioxide. Structures with delocalized electrons always have greater stability than similar structures without delocalized electrons.

Delocalized electrons are not confined between two adjacent bonding atoms, but actually extend over three or more atoms. 10.8

Electron density above and below the plane of the benzene molecule. 10.8

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

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

( ) - bond order = 1 2 Number of electrons in bonding MOs Number of electrons in antibonding MOs ( - ) bond order ½ 1 ½ 10.7

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