Chapter 10. VSEPR Model: Geometries

Transcription

1 Chapter 10 Molecular Geometry VSEPR Model: Geometries Valence Shell Electron Pair Repulsion Theory Electron pairs repel and get as far apart as possible Example: Water Four electron pairs Farthest apart is a tetrahedron 1

2 VSEPR Model: Geometries Experimentally, we only observe the atoms and bonds between them. Water has a bent geometry VSEPR Model: Geometries Start with the Lewis structure Count number of repelling groups around an atom Bonds (single, double, triple) Lone pairs # repelling groups determines their arrangement 2

3 VSEPR Model: Geometries # Repelling Groups Arrangement of Groups Approx. Bond Angles 2 Linear 180 o 3 Trigonal 120 o 4 Tetrahedral 109 o 5 Trigonal Bipyramidal 120 o, 90 o 6 Octahedral 90 o You should know the arrangement and the approximate bond angles. VSEPR Model: Geometries What is the arrangement of groups and the geometry of each molecule below? 3

4 VSEPR Model: Geometries XeF 4 Arrangement: octahedral Geometry: square planar Acetone VSEPR Model: Geometries 4

5 VSEPR Model: Geometries Ethanol VSEPR Model: Bond Angles We can make better estimations of bond angles Some groups repel better than others Repelling power: Lone pair > triple bond > double bond > single bond Compare CH 4, NH 3, and H 2 O Arrangement of groups is tetrahedral in all cases Bond angles differ slightly 5

6 VSEPR Model: Bond Angles Double bonds have more repelling power than single bonds. Lone pairs are even better repellers VSEPR Model: Bond Angles Draw the Lewis structure of each molecule below. Draw a picture of each molecule, showing its geometry and predicting approximate bond angles CH 2 CH 2 O 3 IF 3 6

7 Valence Bond Theory Quantum mechanical theory of bond formation Extension of Lewis theory Lewis: Atoms share a pair of electrons to form bond VB: Orbitals of atoms overlap to form bonds Only orbitals containing unpaired valence electrons Example HF Unpaired valence electrons: H = 1s F = 2p Valence Bond Theory Advantage of VB Theory over Lewis Explains why covalent bond is formed Increased electron density between positive nuclei Nuclei are attracted to electrons between them 7

8 Sigma Bonds These bonds are classified as sigma bonds Rotational symmetry about internuclear axis Hybrid Orbitals Our model of bond generation 2 electrons per bond Only unpaired electrons involved All atoms get full octet of electrons Bonds formed by overlap of atomic orbitals Seems to conflict with carbon 1s 2 2s 2 2p 2 If we use only unpaired e-, how many bonds predicted? How many bonds do we know C always forms? 8

9 Hybrid Orbitals Carbon always forms 4 bonds Explanation Atom forms new orbitals in process of forming bonds 1s 2s 2p 1s sp 3 New orbitals are called hybrid orbitals Orientation of hybrid orbitals is same as VSEPR group arrangements Hybrid Orbitals Assume that hybrid orbitals always are formed before bonding Can predict hybridization of central atom from arrangement of e - pairs Arrangement of e- pairs Hybridization 2 sp 3 sp 2 4 sp 3 5 sp 3 d 6 sp 3 d 2 9

10 Hybrid Orbitals Example: CH 4 (methane) Lewis structure: Four repelling groups Tetrahedral geometry sp 3 hybridization Four hybrid orbitals Four sigma bonds Hybrid Orbitals Predict hybridization of central atoms in following molecules: O in H 2 O Two C s in CH 3 CHO S in SF 6 10

11 Multiple Bonds Example: Ethylene Each C has 3 repelling groups sp 2 hybrid 1s sp 2 2p sp 2 hybrids form σ-bonds with H atoms and other C atom Multiple Bonds What about the remaining unpaired e - in the 2p orbital on each C atom? Parallel p-orbitals overlap and form a π-bond 11

12 Multiple Bonds Single bond: Double bond: Triple bond: one σ one σ and one π one σ and two π How many σ and π bonds in CH 2 =C=CH 2? Check blog. I will add a question about this molecule. Sigma and Pi Bonds Electron density in σ-bonds is concentrated between nuclei Electron density in π-bonds is above and below nuclei Electrons in a π-bond are less tightly held than those in a σ-bond 12

13 Sigma and Pi Bonds A π-bond is weaker than a σ-bond A double bond is stronger than a single bond, but not twice as strong A π-bond is more reactive than a σ-bond Example: Bromination of 2-butene Alkanes (C n H 2n+2 ) Sigma and Pi Bonds Example: n-butane: C 4 H 10 All bonds are sigma: strong and not readily broken Alkanes are not very reactive Used as solvents and fuels Reactivity of organic compounds mainly dependent on multiple bonds and functional groups CH 3 CH 2 -OH CH 3 -NH 2 13

14 Dipole Moments Polar bonds have unequal sharing of electrons The dipole moment is a measure of the polarity High dipole moment = high polarity Not all bonds are polar Dipole Moments Identical atoms: F F Completely nonpolar Atoms with similar electronegativities: C H ΔEN 0.4 Dipole moment is very small Bond is essentially nonpolar 14

15 Molecular Dipole Moment Molecules have dipole moments Resultant of dipole moments of bonds Example: water O H bond: EN difference = = 1.4 Polar Two polar bonds point in same direction Reinforce each other Water molecule is polar (has a significant dipole moment) Molecular Dipole Moments For a molecule to be polar It must have polar bonds Its geometry must not balance If all bonds are nonpolar or essentially nonpolar, then the molecule is nonpolar. CH 3 CH 2 CH 3 C-C bond: Completely nonpolar C-H bond: ΔEN = = 0.4 Essentially nonpolar The molecule is nonpolar 15

16 Molecular Dipole Moment Polar bonds do not necessarily mean a molecule is polar Polar bonds can balance each other Example: CO 2 The dipole moment of CO 2 is zero Molecular Dipole Moment If all bonds in an arrangement (linear, trigonal, tetrahedral, etc) are identical, then the molecule will be nonpolar Example: CF 4 C-F electronegativity difference ΔEN = = 1.5 Polar bond Identical bonds in all positions Dipole moments of bonds will balance Molecule has zero dipole moment 16

17 Molecular Dipole Moment Rank these isomers in order of increasing dipole moment Molecular Dipole Moment Polar bonds are local properties of a molecule Overall polarity of molecule depends on relative significance of polar and nonpolar groups Which of these molecules would be more polar? CH 3 CH 2 CH 2 CH 2 NH 2 CH 3 CH 2 NH 2 17

18 End of material for Exam 4 Molecular Orbital Theory Molecular orbital theory Looks at molecule as a whole All orbitals of atoms in molecule combine to form molecular orbitals MO theory problems from text 52, 54, 55, 56, 58, 63, 64, 65, 68,

19 Interference of Waves Constructive In phase Destructive Out of phase Noninteracting Physically separated Molecular Orbital Theory Atomic orbitals combine to form molecular orbitals Three ways they can combine Antibonding: destructive interference Destabilizes molecule Nonbonding: non-interacting Neither stabilizes or destabilizes molecules Bonding: constructive interference Stabilizes molecule 19

20 Molecular Orbital Theory Atomic orbitals combine to form molecular orbitals # molecular orbitals = # atomic orbitals 2 atomic orbitals => 2 molecular orbitals One bonding, one antibonding 3 atomic orbitals => 3 molecular orbitals One bonding, one nonbonding, one antibonding H 2 molecule MO Theory: H 2 Molecule Combine two 1s atomic orbitals (one from each H atom) Get one bonding and one antibonding molecular orbital Bonding = σ; antibonding = σ* 20

21 MO Theory: H 2 Molecule Feed electrons from atomic orbitals into mo s Each H atom: 1 e - H 2 molecule: 2 e - MO Theory: Bond Order Bond order: indicator of bond strength in molecule Bond order = ½(# bonding e - - # antibonding e - ) BO = 0 Molecule not stable BO = 1 Single bond BO = 2 Double bond Etc. What is bond order for H 2 molecule? 21

22 MO Theory: He 2 Does MO theory predict that He 2 will be stable? What about He 2+? MO Theory: Homonuclear Diatomics There are higher MO energy levels Li 2, B 2, C 2, N 2 O 2, F 2 22

23 MO Theory: Homonuclear Diatomics What is the bond order of the carbide ion, C 2 2-? Use Li 2, B 2, C 2, N 2 energy level diagram Ten electrons Four from each carbon, two extra from -2 charge MO Theory: Homonuclear Diatomics Predict the magnetism of N 2 and O 2. Dia- or paramagnetic? 23

24 MO Theory: Homonuclear Diatomics Make predictions about O 2 molecule and ions Which has the strongest bond: O 2,O 2+, or O 2-? Why Use MO Theory? MO theory is used most of the time to calculate properties of large molecules Predicts energy levels, so can be compared with electronic spectroscopy experiments Better at predicting magnetic properties Better method of describing molecules for which electrons are shared by more than two atoms 24

25 MO Theory: Ozone O 3 : two resonance hybrids Each oxygen is sp 2 hybridized Leaves one extra p-orbital on each oxygen Focus on these p-orbitals MO Theory: Ozone Three p-orbitals Combine to form three molecular orbitals Antibonding Nonbonding Bonding 25

26 MO Theory: Ozone Bonding molecular orbital is most important This is a delocalized molecular orbital Spread over three atoms Contains 2 electrons Delocalized Molecular Orbitals Compounds that must be described by resonance Example: carbonate ion The lowest bonding molecular orbital usually is delocalized 26

27 Delocalized Molecular Orbitals Benzene, C 6 H 6, is most important example Lowest energy bonding molecular orbital is doughnut-shaped, above and below ring Most common symbol for Benzene is below. Molecules with delocalized molecular orbitals generally are more stable than molecules with localized double bonds 27