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14.2 HYBRIDIZATION ESSENTIAL IDEA Hybridization results from the mixing of atomic orbitals to form the same number of new equivalent hybrid orbitals that can have the same mean energy as the contributing atomic orbitals. NATURE OF SCIENCE (2.2) The need to regard theories as uncertain – hybridization in valence bond theory can help explain molecular geometries, but is limited. Quantum mechanics involves several theories explaining the same phenomena, depending on specific requirements. UNDERSTANDING/KEY IDEA 14.2.A A hybrid orbital results from the mixing of different types of atomic orbitals on the same atom. APPLICATION/SKILLS Be able to explain the formation of sp3, sp2 and sp hybrid orbitals in methane, ethene and ethyne. HYBRIDIZATION Definition – model that describes the changes in the atomic orbitals of an atom when it forms a covalent compound. The new set of orbitals are called “Hybrid Orbitals”. What about carbon? Carbon almost always forms four covalent bonds. Its electron configuration though is 1s22s22px12py1 which would lead one to expect carbon to only form two covalent bonds due to the two singly occupied orbitals. Because carbon does form 4 and not 2 covalent bonds, its ground state electron configuration changes during bonding. The process of “excitation” occurs so that one of the “2s” electrons is excited to the unoccupied “p” orbital, thus giving it four singly occupied orbitals available for bonding. Why hybridization? Let’s look at methane – CH4 It has a tetrahedral shape with identical bonds and bond angles of 109.5 degrees. We know that the bonding is happening between the valence electrons of the carbon and four hydrogen atoms. All of the hydrogen valence electrons come from the 1s orbitals. However, the carbon atomic orbitals available for bonding are 1 – s and 3-p orbitals after the process of excitation. It would be expected that the energies from these bonds would be different since they are being formed from different types of orbitals. However, we know that the four C-H bonds in methane have the same energy so somehow the orbitals have been changed and been made equal during the bonding process. When unequal atomic orbitals mix to form new hybrid atomic orbitals which are the same as each other but different from the original orbitals, hybridization has occurred. Comments about hybrid orbitals 1. They do not exist in isolated atoms. 2. They are found only in covalent compounds. 3. They are equivalent in a compound. 4. The number of hybrid orbitals in a bonded atom is equal to the number of atomic orbitals used to form the hybrid orbitals. 5. The type of hybrid orbitals depends upon the electron domain geometry. 6. The atom is able to form stronger covalent bonds using hybrid orbitals. 3 sp Hybridization The previous example which explained the bonding in methane, CH4, is a classic example of sp3 hybridization. This type of hybridization occurs when the three “p” orbitals and one “s” orbital hybridize to form four identical sigma bonds. The shape is tetrahedral and has bond angles of 109.5 degrees. 2 sp Hybridization When carbon forms a double bond as in ethene, it undergoes sp2 hybridization. This type of hybridization occurs when the three “p” orbitals and one “s” orbital hybridize to form three hybrid orbitals and leaves one unhybridized “p” orbital. The shape is trigonal planar with bond angles of 120 degrees. The unhybridized “p” orbitals overlap sideways forming a pi bond. sp Hybridization When carbon forms a triple bond as in ethyne, it undergoes sp hybridization. This type of hybridization occurs when the three “p” orbitals and one “s” orbital hybridize to form two hybrid orbitals and leaves two unhybridized “p” orbitals. The shape is linear with bond angles of 180 degrees. The unhybridized “p” orbitals overlap sideways forming two pi bonds. Ref: jahschem.wikispaces.com APPLICATION/SKILLS Be able to identify and explain the relationships between Lewis structures, electron domains, molecular geometries and types of hybridization. (This is a review of the information found in ppt 4.3) LINEAR A linear molecule has two electron domains. The angle is 180 degrees and it has “sp” hybridization. The Lewis structure has no “lone pairs” of electrons. TRIGONAL PLANAR A trigonal planar molecule has 3 electron domains. It has angles of 120 degrees and “sp2” hybridization. The bent molecule can also have 3 effective pairs if it has one lone pair of electrons. TETRAHEDRAL A tetrahedral molecule has four electron domains. It has angles of 109.5 degrees and “sp3” hybridization. Trigonal pyramidal and “bent” with 2 lone pairs can also have this geometry. TRIGONAL BIPYRAMIDAL A trigonal bipyramidal molecule has 5 electron domains. It has angles of 90 and 120 degrees and “dsp3” hybridization. OCTAHEDRAL An octahedral molecule has 6 electron domains. It has angles of 90 degrees and “d2sp3” hybridization.