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CHEM 1320 Study Guide - Ethane

5 Pages
Fall 2013

Course Code
CHEM 1320

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Chapter 10: Chemical Bonding II
Atomic Orbitals (AOs)
Quantum theory:
Electrons exist only at discrete energy levels
Electrons have wave properties
Wavefunction: probability that an electron will be at a
certain location
AO: Area around a nucleus likely to contain an electron
Each orbital can have 2 electrons (max)
How does Lewis theory explain the bonds in H2 and F2?
Sharing of two electrons between the two atoms
Bond Distance Energy Bond Length Overlap of
H2436.4 kJ/mol 74pm 2 1s
F2150.6 kJ/mol 142pm 2 2p
Valence Bond Theory:
-Valence Bond model:
-Covalent bonds form by combining valence e- orbitals
-AOs combine, overlap, and orbit both nuclei
-Overlap: along an imaginary line connecting the nuclei
Overlap of two “s” electrons: σ-bond
Hybrid Orbitals:
-Methane (CH4)
-4 C-H bonds, all the same
-All sp3-s bondsσ
Hybridization–mixing of two or more atomic orbitals  hybrid
1. Mix at least 2 nonequivalent atomic orbitals (s and p). Hybrid
orbitals: different shape
2. # Hybrid orbitals = # atomic orbitals used
3. Covalent bonds are formed by:
a. Overlap of atomic orbitals
b. Overlap of hybrid orbitals with atomic orbitals
c. Overlap of hybrid orbitals
Ethane (C2H6)
C-H are a sp3-s bondsσ
C-C is a sp3-sp3 bondσ
Lewis theory and NH3:
If the bonds form from overlap of 3 2p orbitals on nitrogen with
the 1s orbital on each hydrogen atom, what would the molecular
geometry of NH3 be?
How do I predict the hybridization of an atom?
Count the # of lone pairs AND the # of atoms bonded
*7A (if they have one bond) and H are unhybridized
# of Lone Pairs
# of Bonded Atoms Hybridization Examples
2 sp BeCl2
3 sp2BF3
4 sp3CH4, NH3, H2O
5 sp3d PCl5
6 sp3d2SF6
3rd Period (and up)

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Chapter 10: Chemical Bonding II 11/13  Atomic Orbitals (AOs) Quantum theory: Electrons exist only at discrete energy levels  Electrons have wave properties Wavefunction: probability that an electron will be at a               certain location AO: Area around a nucleus likely to contain an electron  Each orbital can have 2 electrons (max) How does Lewis theory explain the bonds in H  an2 F ? 2 Sharing of two electrons between the two atoms Bond Distance Energy    Bond Length     Overlap of H 2 436.4 kJ/mol 74pm 2 1s F2 150.6 kJ/mol 142pm 2 2p Valence Bond Theory: ­Valence Bond model: ­ ­Covalent bonds form by combining valence e  orbitals ­AOs combine, overlap, and orbit both nuclei ­Overlap: along an imaginary line connecting the nuclei Overlap of two “s” electrons: σ­bond Hybrid Orbitals: ­Methane (CH ) 4 ­4 C­H 3onds, all the same ­All sp ­s σ bonds Hybridization–mixing of two or more atomic orbitals  ▯hybrid  orbitals 1. Mix at least 2 nonequivalent atomic orbitals (s and p). Hybrid  orbitals: different shape 2. # Hybrid orbitals = # atomic orbitals used 3. Covalent bonds are formed by: a. Overlap of atomic orbitals b. Overlap of hybrid orbitals with atomic orbitals c. Overlap of hybrid orbitals Ethane (C H2) 6 3 C­H are a sp3­s 3 bonds C­C is a sp ­sp  σ bond Lewis theory and NH : 3 If the bonds form from overlap of 3 2p orbitals on nitrogen with  the 1s orbital on each hydrogen atom, what would the molecular  geometry of NH  be3 How do I predict the hybridization of an atom? Count the # of lone pairs AND the # of atoms bonded *7A (if they have one bond) and H are unhybridized # of Lone Pairs + # of Bonded Atoms  Hybridization Examples 2 sp
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