Chapter 10 Theories of Covalent Bonding 1 Atomic Orbitals Molecules
Bonding and 2 Molecular Structure Questions How are molecules held together? Why is O 2 paramagnetic? And how is this property connected with bonding in the molecule? How can noble gases form molecules such as XeF 2, XeF 4, XeOF 4, XeO 3, etc? N 2 with triple bond Phosphorus is a tetrahedron of P 4 atoms.
11.1 Valence Bond (VB) Theory: The Orbital Overlap Model of Bonding by Linus Pauling (1901-1994) The Central Themes of VB Theory A covalent bond forms when the orbitals of two atoms overlap and a pair of electrons with opposite spins occupy the overlap region. 3 The greater the orbital overlap, the stronger the bond. Extent of orbital overlap depends on orbital shape & direction Hydrogen fluoride, HF. Fluorine, F 2.
VB Theory and Orbital Hybridization The orbitals that form when bonding occurs are different from the atomic orbitals in the isolated atoms. If no change occurred, we could not account for the molecular shapes that are observed. Atomic orbitals mix or hybridize when bonding occurs to form hybrid orbitals. The spatial orientation of these hybrid orbitals correspond with observed molecular shapes. 4
Features of Hybrid Orbitals 5 The number of hybrid orbitals formed equals the number of atomic orbitals mixed. The type of hybrid orbitals formed varies with the types of atomic orbitals mixed. The shape and orientation of a hybrid orbital maximizes overlap with the other atom in the bond. Formation and orientation of sp hybrid orbitals and the bonding in BeCl 2. atomic orbitals hybrid orbitals One 2s and one 2p atomic orbital mix to form two sp hybrid orbitals.
Types of Hybridization The 5 types of orbital hybridization correspond to the 5 electron group arrangements in VSEPR theory. Available atomic orbitals mix to form a new set of HYBRID ORBITALS that will give the maximum overlap in the correct geometry 6 # bond shape hybrid # p left 2 linear sp 2 3 trigonal planar sp 2 1 4 tetrahedral sp 3 0 5 trigonal bipyramidal sp 3 d 6 octahedral sp 3 d 2
sp Hybridization Formation and orientation of sp hybrid orbitals and the bonding in BeCl 2. 7 One 2s and one 2p atomic orbital mix to form two sp hybrid orbitals The two sp orbitals are in linear arrangement, 180 apart. Each half-filled sp orbital overlaps with the halffilled 2p orbital of a Cl atom.
Bonding in BF 3 sp 2 Hybridization Bond 8 angle 120 o Overlap of B and F orbitals to form BF 3 - The three sp 2 orbitals point to the corners of an equilateral triangle, their axes 120 apart. - Each half-filled sp 2 orbital overlaps with the half-filled 2p orbital of a F atom. One 2s and two 2p atomic orbitals mix to form three sp 2 hybrid orbitals
sp 3 Hybridization 9 The sp 3 hybrid orbitals in CH 4 One 2s and three 2p atomic orbitals mix to form four sp 3 hybrid orbitals Overlap of C and H orbitals to form CH 4 - The four sp 3 orbitals point to the corners of a tetrahedral shape, their axes 109 apart. - Each half-filled sp 3 orbital overlaps with the half-filled 1s orbital of a H atom.
The sp 3 hybrid orbitals in NH 3 10 The sp 3 hybrid orbitals in H 2 O The N lone pair occupies an sp 3 hybrid orbital, giving a trigonal pyramidal shape. The O lone pairs occupy two sp 3 hybrid orbitals, giving a bent shape
sp 3 d Hybridization The sp 3 d hybrid orbitals in PCl 5. 11 The formation of more than four bonding orbitals requires d orbital involvement in hybridization. One 3s, three 3p, and one 3d atomic orbitals mix to form five sp 3 d hybrid orbitals Overlap of P and Cl orbitals to form PCl 5 - The five sp 3 d orbitals point to the corners of a trigonal bipyramidal shape, their axes 120 and 90 apart. - Each half-filled sp 3 d orbital overlaps with the half-filled 3p orbital of a Cl atom.
sp 3 d 2 Hybridization 12 The sp 3 d 2 hybrid orbitals in SF 6 The formation of more than four bonding orbitals requires d orbital involvement in hybridization. One 3s, three 3p, and two 3d atomic orbitals mix to form six sp 3 d 2 hybrid orbitals Overlap of S and F orbitals to form SF 6 - The six sp 3 d 2 orbitals point to the corners of a square pyramidal shape, their axes 90 apart. - Each half-filled sp 3 d 2 orbital overlaps with the half-filled 2p orbital of an F atom.
From molecular formula to hybrid orbitals 13 Table 11.1 Composition and Orientation of Hybrid Orbitals.
14 Bonding in Glycine sp 3 sp 3 H N H H C H O C sp 2 O H sp 3
Problems 15 Use Valence Bond (VB) Theory to describe the bonding in H 3 O +, CH 3 NH 2, and SF 4. Strategy: Write the Lewis dot structure for the molecule. The electron group geometry around each central atom determines the hybrid orbital set for that atom.
11.2 Modes of Orbital Overlap & 16 the Types of Covalent Bonds A sigma (σ) bond is formed by end-to-end overlap of orbitals. A pi (p) bond is formed by sideways overlap of orbitals. All single bonds are σ bonds. A p bond is weaker than a σ bond because sideways overlap is less effective than end-toend overlap. A double bond consists of one σ bond and one p bond. A triple bond consists of one σ bond and two p bonds
The s bonds in ethane (C 2 H 6 ) 17 Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display. both C are sp 3 hybridized s bond formed by s-sp 3 overlap End-to-end sp 3 -sp 3 overlap to form a s bond A σ bond is cylindrically symmetrical, with its highest electron density along the bond axis. There is relatively even distribution of electron density over all s bonds.
The s and p bonds in ethylene (C 2 H 4 ) 18 Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display. unhybridized 2p orbitals A p bond has two regions of electron density.
The s and p bonds in acetylene (C 2 H 2 ) 19 Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Each C is sp hybridized and has two unhybridized p orbitals.
Electron density and bond order in ethane, ethylene, and acetylene 20 A double bond is less than twice as strong as a single bond, because a p bond is weaker than a σ bond. However, in terms of bond order, a single bond has BO = 1, a double bond has BO = 2, and a triple bond has BO = 3.
Problem 21 Use VB Theory and Orbital Hybridization to describe the bonding in acetone, CH 3 COCH 3, and acetonitrile, CH 3 CN.
Orbital Overlap & Molecular Rotation There is restricted rotation around C=C bond. 22
Consequence of Multiple Bonds The two isomers of 1,2-dichloroethylene (C 2 H 4 Cl 2 ) are different molecules 23 cis-1,2-dichloroethylene polar molecule, bp 60 0 C trans-1,2-dichloroethylene non-polar molecule, bp 47.5 0 C
Limitation of VB Theory: The Paramagnetic Properties of O 2
11.3 Molecular Orbital (MO) Theory & Electron Delocalization The combination of orbitals to form bonds is viewed as the combination of wave functions. Atomic wave functions (AOs) combine to form molecular wave functions (MOs). # of AOs combined always equals # of MOs formed. 25
Key Distinctions between VB & MO Theories 26 Valence Bond (VB) Theory and Orbital Hybridization - closely related to Lewis s idea of bonding and lone pairs of localized electrons on a particular atom. - used for qualitative, visual picture of molecular structure, particularly for polyatomic molecules and all molecules in ground (lowest energy) states. Molecular Orbital (MO) Theory - spread out or delocalized electrons in MOs. - used for quantitative picture of bonding, especially for molecules in excited states. - used in explaining magnetic and spectral properties (e.g., the colors of compounds). - the only theory that can describe accurately the bonding 2009 Brooks/Cole - Cengage in molecules such as O and NO.
Central Themes of MO Theory Addition of AOs forms a bonding MO, which has a region of high electron density between the nuclei. Subtraction of AOs forms an antibonding MO, which has a node, or region of zero electron density, between the nuclei. Bonding MO and Anti-bonding MO have different shapes and energies 27 The bonding MO is lower in energy and the antibonding MO is higher in energy than the AOs that combined to form them. Contours and energies of H 2 bonding s and antibonding s* MOs
Molecular Orbital Diagrams An MO diagram, just like an atomic orbital diagram, shows the relative energy and number of electrons in each MO. The MO diagram also shows the AOs from which each MO is formed. 1. # of MO s = # of atomic orbitals used. 2. Bonding MO is lower in energy than its parent atomic orbitals. Antibonding MO is higher. 3. Atomic orbitals combine to form MOs most effectively when the atomic orbitals are of similar energy.
Filling MOs with Electrons 1. Electrons assigned to MO s of increasing energy 2. An MO can hold a maximum of 2 electrons with opposite spins 3. Orbitals of equal energy are half-filled, with spins parallel, before any of them are filled
Bond Order Bond order = 1/2 [# e- in bonding MOs - # e- in antibonding MOs] Hydrogen (H 2 ) Molecule: (s 1s ) 2 (s* 1s ) 0 BO = ½(2-0) = 1 Dihelium (He 2 ) Molecule (s 1s ) 2 (s* 1s ) 2 BO = ½(2-2) = 0
Learning Check What is the electron configuration of the H 2 + ion in MO terms. Compare the bond order of this ion with He 2 + and H 2-. Do you expect H 2 + to exist? Strategy: Fill the number of valence electrons in the MO diagram or write electron configuration in MO terms. Then calculate Bond Order.
Homomuclear Diatomic Molecules of Period 2 Elements Bonding in s-block homonuclear diatomic molecules Li 2 Li 2 bond order = 1 Be 2 Be 2 bond order = 0
Shapes and energies of s and p MOs from combinations of 2p atomic orbitals
Relative MO energy levels for Period 2 homonuclear diatomic molecules without 2s-2p mixing with 2s-2p mixing MO energy levels for O 2, F 2, and Ne 2 Note the change in energy of s 2p and p 2p MOs with s-p mixing MO energy levels for B 2, C 2, and N 2 s-p mixing occurs when s & p-orbitals are close in energy (< ½ filled 2p, or < 12 valence electrons)
MO occupancy and molecular properties for B 2 through Ne 2. The MO theory fully explains the paramagnetic properties of O 2. Note the change in order of energy of s 2p and p 2p MOs on going from N 2 to O 2.
Learning Check The cation O 2 + and N 2 + are important components of Earth s upper atmosphere. Write electron configurations in MO terms for these ions. Predict their bond orders and magnetic behavoirs. Strategy: Count the number of valence electrons and place them in the MO diagrams.
Electron Configurations for Heteronuclear Diatomic Molecules The MO diagram for HF. - No p MOs. The s MO between 1s of H and 2p of F. As F is much higher in EN than H, it holds electrons tighter resulting the bonding MO is closer F 2p orbitals. - The 2 lone pairs electrons belong to F and do not form bonds. They are in nonbonding MOs.
Electron Configurations for Heteronuclear Diatomic Molecules The MO diagram for NO. - Total 11 valence electrons: s-p mixing and electron configuration: (s 2s ) 2 (s 2s *) 2 (p 2p ) 4 (s 2p ) 2 (p 2p *) 1 - B.O. = 2.5, consistent with 2 resonance structures - the single electron is closer to N than to O
Problems Write electron configurations in MO terms and calculate bond order for CO, and ClF.
Polyatomic Molecules: Resonance and MO Theory Molecular Orbitals in Ozone, O 3, molecule The p electrons delocalize throughout the molecule Fig. 9-22, p. 431
Molecular Orbitals in Benzene, C 6 H 6, molecule Fig. 9-24, p. 432