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Chemical Bonding II Valence Bond & Molecular Orbital Theory. Valence Bond Theory. Orbital Overlap in Chemical Bonds. Bonding Theories - Hybridization.
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Valence Bond Theory Orbital Overlap in Chemical Bonds
Bonding Theories - Hybridization The Valence Shell Electron Pair Repulsion approach works very well for many covalent compounds. It can be used to predict molecular shape, bond angles and molecular polarity. Several bonding theories have been developed to explain how the central atom rearranges its orbitals to minimize electronic repulsion.
Bonding Theories - Hybridization One of the simpler approaches involves hybridization, or the mixing of the atomic orbitals on the central atom. When orbitals hybridize, the properties of the resulting hybridized orbitals are a mixture of the properties of the original orbitals that were mixed together.
Bonding Theories - Hybridization The compounds of carbon will be used for illustration. Methane, CH4, has four equivalent bonds which point towards the corners of a tetrahedron. Theorists have shown that the mixing of the 2s, 2px, 2py, and 2pz orbitals on carbon will produce four equivalent hybrid orbitals that are 109.5o apart.
Bonding Theories - Hybridization Since the hybrid orbitals result from mixing an s orbital with three p orbitals ( the 2px, 2py and 2pz ), the hybrid orbitals that result are called sp3 hybrid orbitals.
Bonding Theories - Hybridization When orbitals are mixed, a hybrid orbital is formed for each atomic orbital that is hybridized. In methane, a total of four atomic orbitals on carbon (2s, 2px, 2py and 2pz) produces four equivalent sp3 hybrid orbitals.
Bonding Theories - Hybridization Four equivalent bonds are formed with hydrogen to form a tetrahedral molecule.
Hybridization of Methane ↓ The hybridization of the orbitals on carbon accounts for the bond angles of the bonds in methane. However, atomic carbon has only two unpaired electrons, and would not be expected to make four bonds. C: 1s2, 2s2, 2p2 =[He] 2s_ 2p _
Hybridization of Methane Scientists proposed that an electron from the 2s level is promoted up to the 2p level.
Hybridization of Methane The carbon atom can now make four bonds instead of two.
Hybridization of Methane Hybridization of the 2s and 2p orbitals produces four equivalent hybrid orbitals.
Hybridization of Methane Although the promotion of an electron and the mixing of orbitals requires energy, there is a net release of energy since four bonds are made, and repulsions are minimized.
Sigma Bonds The bonds between carbon and hydrogen in methane puts electron density on the line connecting the nuclei of the atoms. This type of bond is called a σ (sigma) bond. All covalent molecules contain σ bonds.
sp3 Hybridization Molecules such as water or ammonia are also sp3 hybridized, with hybrid orbitals used for the bonds and any lone pairs of electrons.
Hybridization Trigonal planar geometry results when an s orbital is mixed with two p orbitals. If the molecule is in the xy plane, the px and py orbitals are mixed with an s orbital to form three sp2 hybrid orbitals. The pz orbital remains unchanged, and can be used for making multiple bonds.
Hybridization The energy level diagram shows the mixing of orbitals, with the pz orbital remaining unchanged. ↑ ↑ ↑ ↑ ↑ ↑ ↑ ↑
Hybridization The orbitals on the central atom show trigonal planar geometry, with the pz orbital perpendicular to the molecular plane. z xy plane
Hybridization In ethylene, H2C=CH2, the sp2 orbitals are used to make one of the bonds between the carbon atoms and the bonds between carbon and hydrogen. These bonds have electron density along the internuclear (bond) axis. This type of bond is called a σ (sigma) bond.
Hybridization pz orbitals There are unhybridized pz orbitals on each of the carbon atoms, and they each contain an electron. A second bond forms between the carbon atoms.
The Bond of Ethylene Bonds resulting from side-by-side overlap are called π (pi) bonds. The electrons in the π bond are found above and below the line connecting the two carbon nuclei.
The Bond of Ethylene Even though there are two overlapping lobes (one above the molecular plane and one below it), this is only one π bond, as it involves the sharing of only one pair of electrons.
sp Hybridization Molecules in which the central atom has two atoms attached to it (and no lone pairs of electrons) undergo sp hybridization. Acetylene, HCCH, has a triple bond between the carbon atoms. H-C=C-H
sp Hybridization The shape of the molecule is linear around each carbon atom. The σ bonds between the carbons and between carbon and hydrogen result from the mixing of an s and p orbital on carbon. The result is two sp hybrid orbitals on each carbon.
sp Hybridization If the molecule lies along the x axis, the px orbital on each carbon atom will be used in hybridization. This leaves the py and pz orbitals available for π bonding.
sp Hybridization z x x y
sp Hybridization The σ bond between the carbon atoms arises from overlap of the sp orbitals. The bonds with the hydrogen atoms use the other sp hybrid orbital on each carbon atom.
sp Hybridization There are two p orbitals (the py and pz) on each carbon atom. These are used to make two π bonds between the carbons. ↑ ↑ ↑ ↑ ↑ ↑ ↑ ↑
The Bonding in Acetylene The result is a σ bond along the C-C bond axis, and two π bonds which are perpendicular to each other. pzorbitals pyorbitals
The Bonding in Acetylene The result is a σ bond along the C-C bond axis, and two π bonds which are perpendicular to each other.
sp Hybridization The end result is a carbon-carbon triple bond. H C C H The σ bonds are shown in yellow, and the π bonds are shown in red.
dsp3 Hybridization Central atoms with a total of 5 atoms and lone pairs form trigonal bipyramidal structures.
dsp3 Hybridization The central atom must be in the third period or below, since d orbitals are used to make five dsp3 or (sp3d) hybrid orbitals.
d2sp3 Hybridization Central atoms with a total of 6 atoms and electron pairs form octahedral shapes.
d2sp3 Hybridization The hybridization is d2sp3(or sp3d2), producing six hybrid orbitals that point toward the corners of an octahedron.
Molecular Orbital Theory Valence Bond theory fails to fully explain the bonding in fairly simple molecules. These include molecules or ions with resonance. It also fails to fully explain the bonding of oxygen.
Bonding of Oxygen Oxygen had a double bond which is predicted by the valence bond approach. O=O However, O2is paramagnetic, and contains 2 unpaired electrons. The valence bond approach cannot explain the paramagnetism of oxygen. : : : :
Molecular Orbital Theory Molecular orbital theory views bonds as resulting from the interaction of the wave functions on individual atoms. The waves can interact constructively or destructively. The resulting molecular orbitals belong to the entire molecule, and are not viewed as localized electron pairs.
Molecular Orbital Theory The bonding orbital results in increased electron density between the two nuclei, and is of lower energy than the two separate atomic orbitals.
