Valence shell electrons repel each other Valence shell electrons are arranged geometrically around the central atom to

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1 Molecular Geometry (Valence Shell Electron Pair Repulsion -VSEPR) & Hybridization of Atomic Orbitals (Valance Bond Theory) Chapter 10 Valence Shell Electron Pair Repulsion (VSEPR) Valence shell electrons repel each other Valence shell electrons are arranged geometrically around the central atom to minimize this repulsion Each bond is counted as one pair Each unshared pair of electrons (lone pairs) on the central atom is counted as one pair. Linear (180º) 2 Regions of Electron Density Bond Angles = 180º B-A-B molecules would be non polar A-B-C molecules would be polar 1

2 Trigonal Planer (120º) 3 Regions of Electron Density Lone Pair of Electrons Angular or Bent A-B 3 molecules would be non polar Even if the bonds are polar symmetry in this molecule would cancel the forces yielding the overall molecule non-polar The lone pair of electrons would make this geometry polar and the geometry is no longer symmetric so polar bonds would also yield a polar molecular Electronic vs Molecular Geometries Electronic Geometry Refers to the geometric arrangement of regions of electronic density around a central atom, including lone pairs of electrons Molecular Geometry Describes only the arrangement of atoms around the central atom, excluding lone pairs Tetrahedral (109.5º) 4 Regions of Electron Density Both structures with lone pairs would yield polar molecules A-B 4 molecules would be non polar, even if bonds are polar AB 2 C 2 would yield a polar molecule l Trigonal Pyramidal AB 3 E Angular or Bent AB 2 E 2 A-B 3 C molecules would be polar, especially if A-B or A-C were highly polar Lone pairs of e - occupy more space than bonding pairs. Therefore, the geometric bonding angles tend to be less. 2

3 Electronic vs Molecular Geometries In water, which as 2 lone pairs of electrons and 2 bonds the electronic geometry is tetrahedral the molecular geometry is bent Lone electron pairs take up more space and are held closer to the central atom, therefore the angles seen in molecular geometries with lone pairs are always LESS than that seen in the same geometry with just bonded electrons Trigonal Bipyramidal (90 (90º, 120º, 180º) 5 Regions of Electron Density See-Saw AB 5 would yield a non-polar molecule Linear AB 4 E Polar T-Shaped Non-Polar AB 2 E 3 AB 3 E 2 Polar Lone pair electrons take the equatorial positions, not the axial positions in the trigonal bipyramidal geometry. Why? Octahedral (90º or 180º) 6 Regions of Electron Density AB 5 E AB 6 Non-polar Polar Square Pyramidal AB 4 E 2 Non-polar Square Planar 3

4 VSEPR Steps 1. Draw the Lewis structure for the molecule or ion. 2. Count the total number of regions of electron density (bonding and unshared electron pairs) around the central atom. Double and triple bonds count as ONE REGION OF HIGH ELECTRON DENSITY. An unpaired electron counts as ONE REGION OF HIGH ELECTRON DENSITY. For molecules or ions that have resonance structures, you may use any one of the resonance structures. VSEPR Steps 3. Identify the most stable arrangement of the regions of high electron density (electronic geometry) 4. Determine the positions of the atoms and lone pairs of electrons Lone pairs add to the equatorial positions in the trigonal bipyramidal geometry 5. Identify the molecular geometry based on the positions of the ATOMS (NOT the electron pairs) Valence Bond Theory Describes how & why atoms share electrons to create bonds based on overlap of atomic orbitals (share a common region in space) In isolated atoms, electrons are arranged in orbitals in a way that leads to the lowest total energy for the atom Usually, these atomic orbitals are not in the right geometry or at the correct energy where more than one atom can share the electrons 4

5 Valence Bond Theory When atoms are nearby each other, atoms can combine or rearrange to form NEW orbitals These orbitals result in total lower energy of the atoms These orbitals are called HYBRID orbitals (meaning mixed) The hybrid orbitals have the correct geometry necessary for overlapping with electron orbitals of other atoms sp 3 Hybridization 4 regions of electron density Hybridize Re-arrange Tetrahedral (109.5º) 4 Regions of Electron Density AB 4, AB 3 U, AB 2 U 2 One s, three p s are hybridized to make four sp 3 Lone pairs of e - occupy more space than bonding pairs. Therefore, the geometric bonding angles tend to be less. 5

6 Valence Bond Theory Of course, we label these HYBRID orbitals with odd symbols just to confuse you The labels reflect ect the number and kind of atomic orbitals that combine to create a set of hybrid orbitals Valence Bond Theory Hybrid Orbitals Areas of e - Density Geometry Atomic Orbitals Mixed Label 2 Linear One s, One p sp 3 Trigonal Planar One s, Two p s sp 2 4 Tetrahedral One s, Three p s sp 3 5 Trigonal One s, Three p, One d sp 3 d Bipyramidal 6 Octahedral One s, Three p, Two d s sp 3 d 2 Linear AB 2 2 Regions of Electron Density One s, one p are hybridized to make two sp 2s 2p sp hbid hybrid sp hbid hybrid 2 sp hybrids 6

7 sp Hybridization Mix One s and One p Orbital 2p 2p x, 2p y Be [He] 2s Hybridize Be [He] _ _ sp Trigonal Planer (120º) 3 Regions of Electron Density AB 3 or AB 2U One s, Two p s are hybridized to make three sp 2 Trigonal Planer (120º) 3 Regions of Electron Density AB 3 or AB 2U 7

8 sp 2 Hybridization Mix One s and 2 p Orbitals 2p _ 2p x, C [He] 2s Hybridize C [He] _ _ _ sp 2 sp 2 Hybridization for C 3 regions of electron density One s orbital e- e- Three sp 2 orbitals + one p orbital e - Three p orbitals + e - e - e - e- e- Only use two p orbitals for hybridization 4 electrons available for bonding, but one is in the p atomic orbital Trigonal Bipyramidal (90 (90º, 120º, 180º) 5 Regions of Electron Density AB 5, AB 4 U, AB 3 U 2, AB 2 U 3 One s, three p s & 1 d are hybridized y to make five sp 3 d 8

9 Octahedral (90º or 180º) 6 Regions of Electron Density AB 6, AB 5 U, AB 4 U 2 One s, three p s & 2 d s are hybridized to make six sp 3 d 2 Compounds which have the central atom of a representative ( A ) element, violate the octet rule, if they have sp 3 d 2 hybridization Valence Bond Theory Hybrid Orbitals Areas of e - Density Geometry Atomic Orbitals Mixed Label 2 Linear One s, One p sp 3 Trigonal Planar One s, Two p s sp 2 4 Tetrahedral One s, Three p s sp 3 5 Trigonal One s, Three p, One d sp 3 d Bipyramidal 6 Octahedral One s, Three p, Two d s sp 3 d 2 Polar vs Nonpolar Covalent Bonds Bonds that are partially ionic are polar electrons are NOT equally shared + - (partial positive and negative charges) + In H - F, F has more electron density around it than H. The larger the electronegativity difference ( EN) the more polar the bond the more electronegative atom will have more electron density be more negative 9

10 Dipole Moments Molecules whose centers of positive and negative charge do not coincide, have an asymmetric charge distribution, and are polar. These molecules have a dipole moment. The dipole moment has the symbol. is the product of the distance,d, separating charges of equal magnitude and opposite sign, and the magnitude of the charge, q. 28 Dipole Moments Molecules that have a small separation of charge have a small Molecules that have a large separation of charge have a large For example, HF and HI: - H - F 1.91 Debye units H - I Debye units 29 Molecular Polarity Molecular polarity results from the uneven charge distribution (electron density) between various atoms in a compound which is covalently bonded Polar molecules have a concentration of negative charge on one end of the molecule l and a concentration of positive charge on the other end Have a positive and negative end called a dipole Dipole moment is distance (d) x charge (q) is a measure of the extent of polarity 10

11 Molecular Polarity Geometric Considerations For a molecule to be polar There must be a least one polar bond (compare ENs) or one lone pair of electrons on the central atom The polar bonds must NOT be arranged so their polarities cancel If more than one lone pair on the central atom, they must NOT be arranged so their polarities cancel Molecular Polarity Geometric Considerations Sigma Bonding Single Bonds A sigma ( ) bond results from head on overlap of atomic or hybridized orbitals A sigma bond occurs for all single bonds 11

12 Sigma Bonding Single Bonds 1s 1s 1s 2p 2p 2p Pi Bonding A pi ( ) bond results from the side on overlap of atomic orbitals Pi bonds only occur in the case of double or triple bonds A pi can only occur when a there is also a sigma bond Double bond 1 and 1 Triple bond 1 and 2 Pi Bonding 2p 2p 12

13 Double Bonds A double bond contains 1 and 1 bond For carbon the hybridized orbital is an sp 2 (leaving 1 p orbitals available for the bond) For ethylene, H 2 C=CH 2, each carbon is trigonal planar sp 2 Hybridization Mix One s and 2 p Orbitals _ 2p _ 2p x, C [He] 2s Hybridize C [He] _ _ _ sp 2 For carbon with an sp 2 hybridized orbital, 1 electron remains in a 2 p orbital. The p orbital forms the pi bond. C-C π Bond 13

14 sp 2 Hybridization Ethylene (also called ethene) sp2 Carbon Double Bond One π bond Triple Bonds A triple bond contains 1 and 2 bond For carbon the hybridized orbital is an sp (leaving 2 p orbitals available for the bonds) The geometry of acetylene, HC CH, is linear sp Hybridization Mix One s and One p Orbital _ 2p _ _ 2p x, 2p y C [He] 2s Hybridize C [He] _ _ sp For carbon with an sp hybridized orbital, 2 electrons remains in 2-2p orbitals. The 2 p orbitals forms 2 pi bonds. 14

15 sp Hybridization Acetylene sp Carbon Triple Bond The pi bonds use unhybridized p atomic orbitals to create the bonding (side by side overlap). Problems with Valence Bond Theory VB theory is useful to understand the physical geometry of compounds, but it is difficult to use a predictive tool VB theory would suggest that O 2 is diamagnetic (spin paired electrons which should be mildly repelled by a magnetic field) - O 2 has been proven experimentally to be paramagnetic (electron spins not paired slightly attracted to magnetic fields) opps Molecular Orbital Theory MO theory suggests that atomic orbitals of different atoms combine to create MOLECULAR ORBITALS Electrons in these MOLECULAR ORBITALS belong to the molecule as whole This contrasts to VB theory which suggests that electrons are shared by simple overlap atomic orbitals or hybridized atomic orbitals 15

16 Bonding and Antibonding Orbitals When 2 atomic orbitals are added together a set of lower energy BONDING orbitals are created a set of higher energy ANTI-BONDING orbitals are created Bonding and Antibonding Orbitals Bonding orbitals have most of the electron (negative) density between the 2 positive nuclei Antibonding orbitals have most of the electron density on the opposite side from the region where the bond must be formed H 2 Molecular Orbitals The * indicates that this is an ANTI-BONDING orbital Anti Bonding 1s 1s * E 1s n e r g y 1s Molecular orbitals which result from head-to-head overlap are called sigma ( ) orbitals Bonding 16

17 Molecular Orbital Energy Diagram for H 2 E n e r g y Bond Order Bond Order = (# Bonding e - - # Antibonding e - ) 2 BO = 1 - represents a single bond BO = 2 - represents a double bond BO = 3 - represents a triple bond BO = 0 - represents no bonding (not a stable molecule) Bond Order for H 2 BO = (2 0)/2 = 1 Predicts a single bond of H 2 17

18 Molecular Orbital Energy Diagram for He 2 BO = (2-2)/2 = 0 Molecular Orbital Energy Diagram for He + 2 BO = (2-1)/2 = 1/2 Molecular Orbital Energy Diagram for Li 2 & Be 2 BO = (4-2)/2 = 1 BO = (4-4)/2 = 0 18

19 Molecular Orbital Energy Diagram for N 2 BO = (8-2)/2 = 3 Triple Bond Pi ( ) Molecular Orbitals Atomic Orbitals Molecular Orbitals Molecular Orbitals which from side-by- side overlap of p orbitals are called pi ( ) orbitals. 2py or 2pz 2p y 2p y 2py or 2pz Molecular Orbital Energy Diagram for O 2 BO = (8-4)/2 = 2 Double Bond 19

20 Paramagnetism of O 2 Compounds with unpaired electrons are attracted to a magnetic field This is called paramagnetism MO theory correctly predicts that O 2 is paramagnetic (the last two electrons added are not spin paired) Valance Bond theory does not make this prediction Compare N 2 and O 2 MO Energy Diagrams Hetronuclear Diatomic Molecules Atomic orbitals of different elements will not be at the same energy level The atomic orbitals of the more electronegative element is lower in energy The energy difference between atomic orbitals is roughly proportional to the electronegativity difference between the elements This requires skewing of the standard homonuclear diaatomic MO energy diagram 20

21 HF MO Because there is a large energy difference ( EN = 1.9) the bond is highly polar The sp orbital has characteristic similar to the 2p atomic orbital of F Non bonding orbitals retain the characteristics of the F atomic orbitals Delocalized Oribitals The equivalent of resonance structures in Molecular Orbital theory is called delocalization of electrons across all the bonds Delocalization of Benzene Benzene (C 6 H 6 ) is example of delocalization of electron density 21

22 CO bond length in Carbonate Ion Bond vs Length (Angstroms) CC bond length in Benzene Length in Angstroms Single, Double or Triple Bond C-C C-O C-N N-N 22

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