Chem 1320 1st Edition Lecture 33Outline of Last Lecture I. Molecular Shapes Outline of Current LectureI. VESPRII. Covalent BondingIII. Double BondIV. Triple BondV. Delocalized BondingCurrent Lecture I. VESPR1. Count the number of valence electrons for all atoms2. Draw a Lewis Structure3. Predict the correct geometry by finding electron domain arrangement4. Use the geometry to identify bond angles and net dipole movementII. Covalent BondingThese are pictures you can get from Lewis Structures, atoms share lone pairs of electrons, and so atomic orbital must overlap. Odd electrons are able to participate in bonds. When drawing it out overlap p orbitals where each atom contributes electrons to the bond. You need 4 equivalent bonds to hybridize the orbitals to accommodate the observed bonding. There is an increased bonding capacity where forming bonds releases energy, which stabilizes the molecule.It reduces electron- electron repulsion by the tetrahedral bonding arrangement. This moves the bonding pairs further away from each other. Sp3 hybridization is a characteristic of 4 electron domains around the central atom especially for small atoms. They only hybridize the number of orbitals needed for bonding. There is no advantage of hybridizing. Sp2 hybridization is a characteristic of 3 electron configurations. Single bonds are located along an axis between the nuclei. ∂ is a single bond and these bonds are symmetric about the bonding axis. These notes represent a detailed interpretation of the professor’s lecture. GradeBuddy is best used as a supplement to your own notes, not as a substitute.III. Double BondsThey use the p orbital to increase bonding capacity. Enough hybrid orbitals are formed in order to satisfy the number of ∂ bonds and the extra p electron is then found in the p orbital that is perpendicular to the molecular plane. The ∂ bond is found the ∏ bond orbital. ∏ bond is weaker than the ∂ bond. A double bond is rigid and stronger than a single ∂ bond. This is written as ∂+∏.IV. Triple Bonds2sp orbitals form ∂ bonds. P orbitals are available to increase bonding capacity of the carbons. Porbitals with a single electron are available for bonding. A triple bond is ∂+2∏. V. Delocalized BondingEach c has an extra p electron in an orbital perpendicular to the plane of the molecule. The orbital overlaps to form a delocalized ∏ bond that is spaced over 2 or more atoms. Resonance describes a delocalized
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