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Chapter 11

Chapter 11 chemical bonding

7 Pages
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Department
Chemistry
Course Code
CHMA10H3
Professor
Ann Verner

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Chapter 11: Chemical bonding 11-1 What a bonding theory should do... - Lewis theory, does not help explain why metals conduct electricity, and how semiconductors work - Potential energy- net energy of the interaction of atoms - At very small internuclear spaces, repulsive force exceeds attractive force and the potential energy is + - At intermediate distances, attractive force predominates and the potential energy is - - In molecules, the nuclei vibrates but the avg internuclar space remains constant (bond length) - The potential energy is the negative of the bond-dissociation energy 11-2 introduction to the valence-bond method - region of high e’ probability in H atom  1s orbital - when 2 H come together, the 1s orbital’s in each atom overlap bond is produced b/c of high e’ probability found in overlapped region. - Valence-bond method covalent bond formation when atomic orbital’s overlap - Covalent bond created when half filled orbital’s overlap OR the overlap of one filled orbital with an empty orbital from another atom - The valence bond method gives localized e’ model of bonding: core e’ and lone pair v e’ retain the same orbital location as in the separated atoms, and the charge density of bonding e’ is concentrated at the region of overlap - Energies differs for diff types of orbital’s valence-bond method implies different bond energies for different bonds - 1s-1s overlap in H 2tom produces a grater bond energy than the overlap of 1s-30 in a H-S bond in H 2 11-3 hybridization of atomic orbital’s: - been describing bonded atoms as having the same kind of orbitals as isolated, nonbounded atoms - solution: modify the atomic orbital’s of the bonded atoms - algebraic combination of equations of 2s & 3 2p orbital’s produces a set of 4 identical orbital’s - these 4 orbital’s, are directed in the tetrahedral fashion and have energies that are intermediate between 2s & 2p - hybridization: process of replacing pure atomic orbital’s with formulated atomic orbital’s ( the new orbital’s are called hybrid orbital’s) - hybridization of 1 s and 3 p orbital’s 4 sp hybrid orbital’s - the 3 p orbital`s move down by 1é4 of the energy difference between the s and p orbital and the s orbital movies up by ¾ of that energy diff ( energy conserved) ( pg 427) - hybridization works well with C contain molecules and therefore is used in organic chemistry Bonding in H O an2 NH 3 - SP orbital’s is occupied by one lonepair e’, only the 3 half-filled orbital’s are involved in bonding (NH ) 3 - SP  trigonal-pyramidal SP Hybrid Orbitals: - Boron has 4 orbitals but only 3 V e’ so therefore hybridization of 1 s and 2 2p 2 orbitals while leaving 1p orbital un-hybridized creating SP hybrid orbital and 1 P orbital 2 - Sp  trigonal-planar 2 - # of orbitals conserved ( 3 SP orbitals and 1 P orbital = 4 orbitals) SP Hybrid Orbitals: - Be has 4 orbitals and only 2 Ve’ - 2s &2p orbital are hybridized  2 SP hybrid orbital’s + un-hibirdized 2 2p orbital’s - Sp hybirdization liner SP d & SP d Hybrid orbitals: - Hybridization of 5 and 6 e’ we need to go beyond s and p sub shells into d 3 - Hybridization of s, 3 p and one d orbitals 5 sp d Hybrid orbital’s (PCl ) 5 3 2 - Hybridization of s,3 p and 2 d orbtials 6 sp d hybrid orbital ( SF 6 Hybrid orbitals and the Valence-shell e`- pair repulsion theory: - Hybridization adopted for central atom should produce the same # of hybrid orbitals as there are v-shell e’ groups 11-4 Multiple Covalent bond: - 2 types of diff orbital’s overlap occur when multiple bond are described by the valence-bond method Bonding in C H 2 4 - 1 of the bonds from C results frm the overlap of Sp from each atom - Overlap occurs along the line joining the nuclei of the two atoms - Orbits that overlap in this end-to-end fashion are called(σ) sigma bond - Second bond results from the overlap of un-hybridized p orbital’sthere is regions of high e’ charge density above and below the plane - The bond produced by this side-to-side overlap of 2 parallel orbitals is called pi bond - shape of molecule is determined only by the orbital’s forming σ - rotation of a double bond is restricted - in C-C multiple bonds the σ bond involves more extensive overlap than does bonddouble bond ( σ) is stronger than a single bond (σ)but not twice as strong. - In a triple bond one is a σ bond and two are bonds. 11-5 - VSEPR theory & valence bond method do not provide explanation of +he electronic spectra of molecules, why oxygen is paramagnetic, or why H 2 stable - molecular orbital theory explains theory assigns e’ in a molecule to a series of orbitals that belong to the molecule as a whole  called molecular orbitals - like atomic orbitals, molecular orbitals are: mathematical functions that relate to finding the probability of an e’ in a certain region - like atomic orbitals, molecular orbitals can only have 2 e’ per orbital with opposing spins - H atoms merge to form a chemical bond: the 2 1s orbital combine by interfering constructively are destructively - Constructive interference: adding the 2 mathematical equations, + sign puts waves in phase - Constructive interference of the 2 wave functions leads to a greater probability of finding e’ between the 2 nuclei - E’ charge density = e’ probability - Result of constructive interference is bonding molecular orbital’s b/c it
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