chapter 9 molecular shapes -shape of molecule is based on bond angles valence shell electron pair...
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Chapter 9Molecular Shapes
-shape of molecule is based on bond angles
Valence Shell Electron Pair Repulsion (VSEPR)
-based on the idea that electron groups repel one another
-for molecules having one central atom, shape is based on:
e- domain- region where e- are likely to be found
one e- domain = a lone pair, single, double or triple bond
-must have the structural formula drawn
-use the following formula
ABnEn
A= central atom
B= # of bonding domains on central atom
E= # of lone pairs of e- on central atom
e- geometry- shows arrangement of e- domains
molecular geometry- arrangement of the atoms
Molecular Shape and Polarity
-to determine if an entire molecule is polar you must look at the bonds and the shape
dipole moment- a measure of the separation of + and – charges in a molecule
bond dipole- dipole moment due only to the two atoms in that bond
-bond dipoles and dipole moments are vector quantities
-have both magnitude and direction and must look at both when determining polarity
-vectors will be added
To determine if a molecule is polar:
-determine the shape
-determine whether the molecule contains polar bonds
-determine whether the polar bonds add together to form a net dipole moment
-if they sum to zero then the molecule is
nonpolar
-if not then the molecule is polar
-if the molecule has lone pairs, it is polar
Examples:a) BrCℓ
polarb) SO2
polar *has lone pairc) SF6
nonpolard) NF3
polar *lone paire) BCℓ3
nonpolar
Valence Bond Theory
-valence e- of atoms in a molecule reside in atomic orbitals which can be s, p, d, f or some hybrid of these
-a chemical bond results from the overlap of two half-filled orbitals, or less commonly from the overlap of a completely filled orbital with an empty orbital
-the shape of the molecule is determined by the geometry of the overlapping orbitals
-overlap of half-filled orbitals does not explain bonding in all types of molecules
-for example: CH4
-b/c there are only two half-filled orbitals, we would assume that C bonds with only two H
-we know it bonds with 4 H
-we must consider hybrid orbitals here
hybridization- standard atomic orbitals are combined to form hybrid orbitals that correspond more closely to the actual distribution of e- in chemical bonds
General Statements Regarding Hybridization
-# of standard atomic orbitals added together = # of hybrid orbitals formed
-the combo of orbitals determines the shapes and energies of orbitals formed
-the specific type of hybridization that forms for a molecule is the one that yields the lowest overall energy for the molecule
-a single bond contains a sigma bond (σ)- head on overlap
-a double bond is made up of one sigma and one pi bond (π)- sideways overlap
-a triple bond is made up of one sigma and two pi bonds
sp3 hybridization
-occurs when there are 4 e- groups (tetrahedral)
-for CH4, one s orbital and three p orbitals hybridize to form four sp3 orbitals of equal energy
-C has four half-filled orbitals to overlap with four half-filled H 1s orbitals
*all are single bonds sigma bonds (σ)
sp2 hybridization and double bonds
-hybridization of one s orbital and two p orbitals results in three sp2 hybrids and one leftover unhybridized p orbital
-occurs when you have 3 e- groups (tri. planar)
-because C is the central atom it undergoes hybridization
*made up of two sigma bonds from the C-H single bonds, and one sigma and one pi bond (π) from the double bond of C=O
**Try H2C2Cℓ2
sp hybridization and triple bonds
-one s orbital and one p orbital hybridize to form two sp orbitals of equal energy
-two p orbitals remain unhybridized
-occurs when there are two e- groups (linear)
Practice:
Tell the type of hybridization of the C atom(s), what types of bonds, and what kinds of orbitals for each.
1) H3C2OH 2) HCN 3) C2H4
1) 1st C= sp3, 4 σ bonds, 1s and 3p orbitals
2nd C= sp2, 3 σ and 1 π bond, 1s, 2p, and 1 unhybridized p orbital
2) sp, 2 σ bonds, 2 π bond, 1s, 1p and 2 unhybridized p orbitals
3) **both interior C the same= sp2, 3 σ bonds, 1 π bond, 1s, 2p, and 1 unhybridized p orbital
localization of e-: σ and π e- are associated only with the two atoms that form the bond
delocalization of e-: e- in π bonds extend over more than just the two atoms
*occurs when a molecule has at least two resonance structures
Molecular Orbital Theory
-orbitals are treated as overlapping the entire molecule, not as “belonging” to individual atoms
-when atomic orbitals are combined to form molecular orbitals, the total number of orbitals does not change
-any time you bond two atomic orbitals you get two molecular orbitals:
bonding orbital (σ or π ) - lower in energy
anti-bonding (σ* or π* ) - higher in energy
*b/c lower energy bonding fills first
*each orbital can hold two e-
*will show if a single, double or triple bond is formed by calculating the bond order
bond order = # e- in bonding - # e- in antibonding
2
-if bond order is zero or less, no bond will form
-σ are one box b/c of head on overlap
-π are two boxes b/c of sideways overlap
s orbitals = σ and σ* for the bonding and antibonding
p orbitals = π, σ, π* and σ* for bonding and antibonding