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Orbital Hybridization & Resonance

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Orbital Hybridization & Resonance Valence Bond Theory Overlap and Bonding • Lewis taught us to think of covalent bonds forming through the sharing of electrons by Atomic adjacent atoms. • In such an approach this can only occur when Orbitals are orbitals on the two atoms overlap according to bad… Walter Heitler and Fritz London . mmmkay Linus Carl Pauling (1901-1994) Nobel prizes: 1954, 1962 Overlap and Bonding Localized Electron Model • Increased overlap From this description of covalent bonding rises brings the electrons the Localized Electron Model: and nuclei closer Orbitals overlap to form a bond between two together while atoms due to simultaneous attractions to both simultaneously nuclei. decreasing electron- Two electrons, of opposite spin, can be electron repulsion. accommodated in the overlapping orbitals (usually one electron is supplied by each of the two • However, if atoms get bonded atoms) too close, the Because of orbital overlap, the bonding electrons internuclear repulsion have a higher probability of being found with greatly raises the along the internucleus axis. energy. This simple picture of orbital overlap So then, how can we account for works well to describe the energies o four bonds to carbon when the associated with the formation of covalent 109 109 o bonds but faces a challenge when you try atomic orbitals for carbon will to explain the geometry and number of only allow for the sharing of bonds expected for multiple atoms. two electron pairs according to Localized Electron Model. For example, CH 4 Atomic should have a tetrahedral Also, how can the tetrahedral orbitals of geometry with four bonded shape that is found C: hydrogens surrounding the experimentally be attained? carbon atom; However, carbon only has two unpaired 2p __ __ __ electrons to use in Hybridize the Orbitals! overlapping orbitals, not 4. 2s _____ 1 In 1930, Pauling described Orbital Hybridization to explain Hybridization the problems that arise between 109 o • Process that changes properties of the predicted electron valence electrons by mixing atomic arrangement (L.E. model) and orbitals to form special orbitals for the predicted molecular geometry bonding (VSEPR theory) 4 C atom orbitals Hybridized Orbitals of hybridize to form C: four equivalent __ __ __ __ sp 3 hybrid atomic Hybridized 3 atomic orbitals. sp hybrid orbitals orbitals orbitals Principles of Hybridization Orbital Hybridization 1. Conservation of orbitals 2. Hybrid orbitals correlate with molecular geometry 3. Energy level of hybrid orbitals is between that of AO’s 4. All bonded atoms hybridize and attain the lowest energy arrangement possible All hybrid orbitals of an atom are said to be DEGENERATE (of equal energy) • Mixing the s and p orbitals yields two degenerate Hybrid Orbitals orbitals that are hybrids of the two orbitals. These sp hybrid orbitals have two lobes like a p orbital. • Consider beryllium: One of the lobes is larger and more rounded as is the s In its ground electronic orbital. state, it would not be able to form bonds because it has no singly-occupied orbitals. • These two degenerate orbitals would align themselves But if it absorbs the small 180 °°° from each other. amount of energy needed • This is consistent with the observed geometry of to promote an electron beryllium compounds: from the 2 s to the 2 p orbital, it can form two linear bonds. 2 Hybrid Orbitals sp hybridization linear species • With hybrid orbitals the orbital diagram for beryllium would look like this. • The sp orbitals are higher in energy than the 1 s orbital but lower than the 2 p. Using a similar model for boron leads to… sp 2 hybridization trigonal planar species …three degenerate sp 2 orbitals. With carbon we get… sp 3 hybridization tetrahedral species …four degenerate sp 3 orbitals. 3 1515 OrbitalOrbital HybridizationHybridization Orbital Hybridization 1. Identify and draw the hybridization for ammonia, NH 3. BONDSBONDS SHAPESHAPE HYBRIDHYBRID REMAINREMAIN 22 linearlinear spsp 22 pp’’ss 33 trigonaltrigonal spsp 22 11 p p planarplanar 4 tetrahedral sp 33 none 4 tetrahedral sp none Notice all bonds lie on the internuclear axis for each orbital. 2. Identify and draw the hybridization Valence Bond Theory for formaldehyde, CH O. 2 • With the inclusion of Resonance and Orbital Hybridization by Pauling, Localized Electron Model grew into the modern Valence Bond Theory. Hybridization is a major player in this approach to bonding describing two ways orbitals can overlap to form bonds between Where is the second bond? atoms. Remember, a bond is an overlap of electron densities between atoms Sigma ( σσσ) Bonds Pi ( πππ) Bonds • Pi bonds are characterized by Side-to-side overlap. Electron density • Sigma bonds are characterized by above and below the internuclear Head-to-head overlap. axis. Cylindrical symmetry of electron density about the internuclear axis. 4 Single Bonds Multiple Bonds Single bonds are always σσσ bonds, because σσσ In a multiple bond one of the bonds is a σσσ overlap is greater, resulting in a stronger bond and the rest are πππ bonds. bond and more energy lowering. Multiple Bonds Multiple Bonds • In a molecule like formaldehyde (shown In triple bonds, as at left) an sp 2 orbital on in acetylene, two carbon overlaps in σσσ sp orbitals form a σσσ fashion with the bond between the corresponding orbital carbons, and two on the oxygen. pairs of p orbitals πππ • The unhybridized p overlap in orbitals overlap in πππ fashion to form the πππ fashion. two bonds. 3.Draw the Lewis structure for ethene Consequences of Multiple Bonding (C 2H4), describe its hybridization and There is restricted rotation around C=C bond. identify each bond in the molecule. This gives rise to cis/trans, or geometric, isomerism for alkenes. 5 - σσσ Investigate NO 3 and πππ Bonding in C2H2 In nitrogen trioxide, the electrons of the double bond are delocalized across the molecule. Due to the morphing of atomic orbitals to sp 2 hybridized orbitals, the H– C ≡ C –H delocalized electrons stabilize the molecule. Therefore, the actual structure should be representative of all three Lewis structures. Delocalized Electrons: Resonance Delocalized Electrons: Resonance When writing Lewis structures for species • In reality, each of the four like the nitrate ion, we draw resonance atoms in the nitrate ion has structures to more accurately reflect the a p orbital. structure of the molecule or ion. • The p orbitals on all three oxygens overlap with the p orbital on the central nitrogen. Delocalized Electrons: Resonance Resonance This means the πππ electrons The organic are not localized between the molecule benzene nitrogen and one of the has six σσσ bonds and oxygens, but rather are a p orbital on each delocalized throughout the carbon atom. ion. 6 The Valence Bond theory, by its nature, does not Resonance account for resonance structures due to the delocalization of electrons, however, it does predict its existence for some molecules. πππ • In reality the electrons in benzene are not But… localized, but delocalized. π It does provide a simple model to describe a • The even distribution of the electrons in visual picture of molecular structure through benzene makes the molecule unusually stable. the use of: • Lewis Structures • VSEPR Theory • Orbital Hybridization • Resonance • Don’t forget bond polarity and any resulting molecular polarity according to molecular geometry. 7.
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