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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 Model(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 The bonded atoms and the lone pairs around a central atom A arrange themselves so as to minimize the electrostatic repulsions among the electron pairs around A. Molecular Geometry: arrangement of Atoms attached to A alone (exclude lone pairs).
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linear linear trigonal planar trigonal planar tetrahedral tetrahedral
B linear trigonal planar trigonal planar tetrahedral tetrahedral Trigonal bipyramidal Trigonal bipyramidal Octahedral Octahedral
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0 lone pairs on central atom
Cl Be 2 atoms bonded to central atom 10.1
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When there are no lone pairs on A, the arrangement of electron pairs (attached atoms & lone pairs) and the molecular geometry (attached atoms alone) are identical. Presence of 1 or more lone pairs on the central atom, A, will distort the molecular geometry. This is because, lone pair-lone pair repulsion > lone pair-bond pair repulsion > bond pair-bond pair repulsion Therefore, the molecular geometry (attached atoms alone) will not be identical to the arrangement of electron pairs (attached atoms & lone pairs) when there are 1 or more lone pairs on the central atom, A. Lone pairs are represented by the letter E
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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 AB2E2 2 2 bent H O
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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 bipyramidal trigonal bipyramidal AB5 5 trigonal bipyramidal seesaw 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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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 SO2 and SF4? S F S O AB4E AB2E seesaw bent 10.1
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Predicting Molecular Polarity
A molecule with polar bonds will be nonpolar if bond dipoles cancel. A molecule will be polar if it has at least 1 polar bond and bond dipoles do not cancel. Generalizations: Molecules with linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral & square planar MOLECULAR geometries will be NONPOLAR if END ATOMS ARE IDENTICAL (bond dipoles cancel). Example: O=C=O is nonpolar but H-C≡N is polar. Although both are linear, in O=C=O, the bond dipoles cancel but in H-C≡N they do not. Molecules with bent, pyramidal, t-shaped, see-saw MOLECULAR geometries are polar (bond dipoles do not cancel). Example: a. CO2 is nonpolar (linear geometry) but SO2 is polar (bent geometry); BF3 is nonpolar (trigonal planar geometry) but NH3 is polar (pyramidal geometry); CCl4 is nonpolar (tetrahedral geometry) but SF4 is polar (seesaw geometry).
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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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Does BF3 have a dipole moment?
No, the bond dipoles cancel: trigonal planar molecular geometry. 10.2
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Does CH2Cl2 have a dipole moment?
Yes, although the molecular geo is tetrahedral, the bond dipoles do not cancel as the end atoms are not the same: 2 are Cl and 2 are H. 10.2
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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 unpaired e- from overlapping atomic orbitals. 10.3
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Change in Potential Energy of Two Hydrogen Atoms
as a Function of Their Distance of Separation As the 2 1S orbitals approach each other, the single electron in each one experiences a simultaneous attraction to both nuclei & hence the potential energy drops. As the distance between the nuclei gets even smaller, they repel and the energy starts to increase. The distance corresponding to minimum potential energy = bond length. 10.3
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In the context of valence bond theory, hybridization is introduced to explain the geometries of molecules. Hybridization is the mixing of 2 or more VALENCE orbitals on an atom to create AN EQUAL NUMBER of DEGENERATE (having the same energy) HYBRID ORBITALS. The type of hybridization is determined by the sum of the lone pairs + number of atoms attached to the atom in question. For instance, the hybridization at C, N & O are identical in CH4, NH3, H2O because, in all of these cases, lone pairs + attached atoms = 4. CH4: 0 lone pairs + 4 attached atoms = 4 NH3: 1 lone pair + 3 attached atoms = 4 H2O: 2 lone pairs + 2 attached atoms = 4
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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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Formation of sp3 Hybrid Orbitals
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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?
Draw the Lewis structure of the molecule. 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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A single bond consists of an end-on-end overlap of pur or hybrid orbitals. Such an overlap is called a sigma overlap. A double bond consists of 1 sigma overlap and 1 parallel overlap. A parallel overlap (example: velcro) is called a pi overlap or pi bond. A triple bond consists of 1 sigma overlap and 2 pi overlaps. Single bond = 1 sigma bond Double bond = 1 sigma bond + 1 pi bond Triple bond = 1 sigma bond + 2 pi bonds.
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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 – The atomic orbitals of all atoms in a molecule mix to form molecular orbitals. Molecular orbitals are delocalized over the entire molecule. Atomic orbitals of similar energy on combining atoms mix to form an equal number of molecular orbitals. 10.6
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We will only consider 1st period & 2nd period diatomics: H2, He2, Li2, Be2, B2, C2, N2, O2, F2, Ne2, and combinations of these (CO, NO etc). The atomic orbitals combine pairwise to form pairs of molecular orbitals which differ in energy. One of the pair will have lower energy than both combining orbitals: bonding molecular orbital. The other will have higher energy than both combining orbitals: antibonding molecular orbital. When atomic orbitals combine in an end–on-end manner sigma molecular orbitals are formed. When they combine in a parallel manner, pi molecular orbitals result. Any pair of s orbitals combines to form 2 sigma molecular orbitals (bonding & antibonding). A pair of p orbitals can either form sigma Mos or pi Mos depending on whether they combine end-on-end or parallel.
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*: antibonding Mos Ground state molecular orbital configuration: lowest energy arrangement of MOs Rules in filling MO’s: Pauli exclusion principle (only 2 electrons in each MO) Aufbau principle: Start filling from the lowest energy MO Hund’s Rule: Fill degenerate orbitals (The two pi bonding MOs & the two pi* antibonding Mos) singly before pairing. Count ALL electrons NOT JUST VALENCE when writing MO configuration.
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Diamagnetic diatomics: Have no unpaired electrons in their MO configuration.
Paramagnetic diatomics: Have unpaired electrons in their MO configuration.
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Bond Order = (number of Bonding electrons – number of antibonding electrons)/2
Larger the bond order, stronger the bond, shorter the bond Single bond: bond order = 1 Double Bond: Bond order = 2 Triple Bond: Bond order =3 No bond = bond order = 0 MO theory allows fractional bond orders (0.5, 1.5, 2.5 etc)
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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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Two Possible Interactions Between Two Equivalent p Orbitals
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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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10.7
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