1 Molecular Geometry and Bonding Theories Brown, LeMay Ch 9 AP Chemistry 2 9.1 – 9.2: V.S.E.P.R. Valence-shell electron-pair repulsion theory Because e - pairs repel, molecular shape adjusts so the valence e - pairs are as far apart as possible around the central atom. Electron domains: areas of valence e - density around the central atom; result in different molecular shapes Includes bonding e - pairs and nonbonding e - pairs A single, double, or triple bond counts as one domain Summary of L m AB n (Tables 9.1 - 9.3): L = lone or non-bonding pairs A = central atom B = bonded atoms Bond angles notation used here: < xº means ~2-3º less than predicted << xº means ~4-6º less than predicted 3 Predicting Molecular Geometries: 1. Sketch the Lewis Structure of the molecule or ion 2. Count total number of electron domains on central atom: a. Bonding pair of electrons (between two atoms) = 1 electron domain b Nonbonding pair of electrons (lone pair) = 1 electron domain b. Nonbonding pair of electrons (lone pair) = 1 electron domain c. Double bond (between two atoms) = 1 electron domain d. Triple bond (between two atoms) = 1 electron domain 3. Use the arrangement of the bonded atoms to determine the molecular geometry.
Transcript
1
Molecular Geometry and
Bonding Theories
Brown, LeMay Ch 9AP Chemistry
29.1 – 9.2: V.S.E.P.R.Valence-shell electron-pair repulsion theory
Because e- pairs repel, molecular shape adjusts so the valence e- pairs are as far apart as possible around the central atom.
Electron domains: areas of valence e- density around the central atom; result in different molecular shapes Includes bonding e- pairs and nonbonding e- pairsg p g p A single, double, or triple bond counts as one domain
Summary of LmABn (Tables 9.1 - 9.3):L = lone or non-bonding pairsA = central atomB = bonded atoms
Bond angles notation used here:< xº means ~2-3º less than predicted<< xº means ~4-6º less than predicted
3Predicting Molecular Geometries:
1. Sketch the Lewis Structure of the molecule or ion 2. Count total number of electron domains on central
atom:a. Bonding pair of electrons (between two atoms) = 1 electron domain b Nonbonding pair of electrons (lone pair) = 1 electron domainb. Nonbonding pair of electrons (lone pair) = 1 electron domainc. Double bond (between two atoms) = 1 electron domaind. Triple bond (between two atoms) = 1 electron domain
3. Use the arrangement of the bonded atoms to determine the molecular geometry.
2
4Tables 9.1 - 9.3
# of e-
domains&
# and type of hybrid orbitals
e- domain geometry
Formula &
Molecular geometry
Predicted bond
angle(s)
Example(Lewis
structure with molecular
shape)
2
Two sp hybrid orbitals
LinearAB2
Linear
180º BeF2
CO2
A|X
X
A|B
B
3
120º
BF3
Cl-C-Cl120ºA
XX
A|B
BB
Three sp2
hybrid orbitals
Trigonal planar
AB3
Trigonal planar
120º
Cl2CO
LAB2
Bent
< 120º
NO21-
A|X
A|B
:B
6Example: CH4
Molecular shape = tetrahedral
Bond angle = 109.5º
H|
H—C—H|HH
109.5º
109.5º109.5º
109.5º
3
4
AB4
Tetrahedral
109.5º
CH4
< 109.5º
XA
X
XX
BA
B
BB
:A
B
Four sp3
hybrid orbitals
or
Tetrahedral
LAB3
Trigonal pyramidal
Ex: NH3 = 107º
NH3
L2AB2 Bent
<<109.5º
Ex: H2O = 104.5º
H2O
X
A
X
XX
BB
:A
B
B:
8Effect of Nonbonding Electrons and Multiple Bonds on Bond Angles
Nonbonding electron pairs exert greater repulsive forces on adjacent electron domains therefore compressing the bond angles
Multiple bonds exert a greater repulsive force on adjacent electron domains therefore compressing the bond angle
9Molecules with Expanded Valence Shells More than the “perfect octet” on the central atom
Electron domains can point in two geometrically distinct positions”
169.3: Molecular Polarity A molecule is polar if its centers of (+) and (-)
charge do not coincide. A bond’s polarity is determined by the difference of
electronegativity between atoms in bond. Partial (+) and partial (-) charges on atoms in a polar bond
can be represented as + and -can be represented as and .
Bond polarity is most often represented by an arrow that points toward the - (most EN atom), showing the shift in e- density. H-Cl:
::
H-Cl:
::
+ -
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The dipole moment () is a vector (i.e., has a specific direction) measuring the polarity of a bond which contains partial charges (Q) that are separated by a distance (r).
= Q r
The sum of the bond dipole moments in a molecule determines the overall polarity of the molecule.1. Draw the true molecular geometry.2. Draw each bond dipole as an arrow (not lone pairs)3. Add the vectors, and draw the overall dipole moment. If
none, then = 0.
Ex: Draw molecular geometries, bond dipole moments, and overall dipole moments. Also, name the e- domain geometry and the molecular geometry.
CO2 BF3 H2O
CCl4 NH3 PH3
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199.4: Covalent Bonding and Orbital Overlap
Valence-bond theory: overlap of orbitals between atoms results in a shared valence e- pair (i.e., bonding pair)
Energy
Figure 9.13:Formation of bond in H2
a. As 2 H atoms approach, the 2 valence e-in the 1s orbitals begin to overlap, becoming more stable.
b. As H-H distance approaches 0.74 Å, energy lowers b/c of electrostaticEnergy
(kJ/mol)0
-436
0.74 ÅH-H distance
energy lowers b/c of electrostatic attraction between the nuclei & the incoming e-.
c. When H-H distance = 0.74 Å, energy is at its lowest because electrostatic attractions& repulsions are balanced. (This is the actual H-H bond distance.
d. When H-H distance < 0.74 Å, energy increases b/c of electrostatic repulsion between 2 nuclei & between the 2 e-.
ab
c
d
209.5: Hybrid Orbital Theory Explains the relationship between overlapping
orbitals (valence bond theory) and observed molecular geometries (VSEPR theory).
21“sp” hybrid orbitals BeF2 (g): observed as a linear molecule with 2 equal-length
Be-F bonds. Valence bond theory predicts that each bond is an overlap of one Be 2s e- and one 2p e- of F. However, Be’s 2s e- are already paired. So…
To form 2 equal bonds with 2 F atoms:1. In Be, one 2s e- is promoted to an empty 2p orbital.
2. The occupied s and p orbitals are hybridized (“mixed”), producing two equivalent “sp” orbitals.
3. As the two “sp” hybrid orbitals of Be overlap with two p orbitals of F, stronger bonds result than would be expected from a normal Be s and F p overlap. (This makes up for energy needed to promote the Be e- originally.)
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Be (ground state) → Be (promoted) → Be (sp hybrid)
2p
2s
Ener
gy →
O bit l “ h ”F F
Orbital “shapes”One s + one p → Two sp orbitals
(to bond with 2 F’s)
A central atom in a Lewis structure with exactly 2 e- domains has sp hybrid orbitals.
23“sp” hybrid orbitals
24“sp2” hybrid orbitals BF3 (g): observed as trigonal planar molecule with 3 equal-
length B-F bonds. However, 2 valence e- in B are paired, and are the s and p e- not at the observed 120º angle.
2pB (ground) → B (promoted) → B (sp2 hybrid)
2s
One s + two p → Three sp2 orbitals(to bond with 3 F’s)
A central atom with exactly 3 e- domains has sp2 hybrid orbitals.
F F F
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25“sp2” hybrid orbitals
26“sp3” hybrid orbitals CH4 (g): observed as tetrahedral
2p
2s
C (ground) → C (promoted) → C (sp3 hybrid)
One s + three p → Four sp3 orbitals(to bond with 4 H’s)
A central atom with exactly 4 e- domains has sp3 hybrid orbitals.
H H H H
27“sp3” hybrid orbitals
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28“sp3d” hybrid orbitals (or dsp3) PCl5 (g): observed as trigonal bipyramidal; forms 5 bonds of equal energy
(* but not equal length: equatorial are slightly longer)
3d
3p
P (ground) → P (promoted) → P (sp3d hybrid)
3s
One s + three p + one d → Five sp3d orbitals(to bond with 5 Cl’s)
A central atom with exactly 5 e- domains has sp3d hybrid orbitals.
Cl Cl Cl Cl Cl
29“sp3d2” hybrid orbitals (or d2sp3) SF6 (g): observed as octahedral; forms 6 equal-length bonds
One s + three p + two d → Six sp3d2 orbitals
A central atom with exactly 6 e- domains has sp3d2 hybrids.
30Non-bonding e- pairs Lone pairs occupy hybrid orbitals, too
Ex: H2O (g): observed as bent; but e- domain is tetrahedral
2p
O (ground) → O (sp3 hybrid)2 bonding pairs
2s
Four sp3 orbitals (2 bonding, 2 non-bonding)
2 non-bonding pairs(lone pairs) H H
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31To predict hybrid orbitals:
1. Draw Lewis structure for the molecule or ion2. Determine the electron domain geometry using the
VSERP modelS if th h b id bit l d d t d t3. Specify the hybrid orbitals needed to accommodate the electron pairs based on their electron geometric arrangement (Table 9.4 pg 361)
329.6: Multiple BondsDraw Lewis structures. For C’s: label hybridization, molecular geometry,
and unique bond angles
C2H6
C2H4
C2H2
C6H6
33Sigma and Pi bondsSigma () bond: Covalent bond that results from axial overlap of orbitals between
atoms in a molecule Lie directly on internuclear axis (line connecting the nuclei) “Single” bonds
Ex: F2
Pi () bond: Covalent bond that results from side-by-side overlap of orbitals
between atoms in a molecule. Are “above & below” and “left & right” of the internuclear axis and
therefore have less total orbital overlap, so they are weaker than bonds
Make up the 2nd and 3rd bonds in double & triple bonds.Ex: O2 N2
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34Sigma () bonds in C2H4
Ex: ethene; C-C -bonds and C-H -bonds result from axial overlap of H s-orbitals and C sp2-orbitals
35Pi () bonds in C2H4
2pC
2s sp2 hybrids bond axially
p orbital bonds side-by-side = bond
Each C has 4 valence e-: 3 e- for 3 bonds 1 e- for 1 bond, which results from side-by-side
overlap of one non-hybridized p-orbital from each C
y= bonds
36Sigma () bonds in C2H2
Ex: ethyne (a.k.a. acetylene) C-C -bond and C-H -bonds result from axial overlap of H s-orbitals and C sp-orbital
13
37Pi () bonds in C2H2
sp hybrids bond axially = bonds
2pC
2s
p orbital bonds side-by-side = bonds
Each C has 4 valence e-: 2 e- for 2 bonds 2 e- for 2 bonds, which result from side-by-side overlap
of two non-hybridized p-orbitals from each carbon
38Sigma () bonds in C6H6
Ex: benzene; C-C -bonds and C-H -bonds result from axial overlap of H s-orbitals and C sp2-orbitals
39Localized vs. Delocalized Bonds
(localized) (delocalized –MINIMUM OF 4 c’S)
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40Delocalized bonds in C6H6
C-C -bonds result from overlap of one non-hybridized p-orbitals from each C
Delocalization of e- in -bonds results in a “double-donut” shaped e- cloud above and below the molecular carbon plane.
419.7: Molecular Orbital (MO) theory
So far we have used valence-bond theory (covalent bonds form from overlapping orbitals between atoms) with hybrid orbital theory and VSEPR theory to connect Lewis structures to observed molecular geometries.MO theory is similar to atomic orbital (AO) theory (s MO theory is similar to atomic orbital (AO) theory (s, p, d, f orbitals) and helps to further explain some observed phenomena, like unpredicted magnetic properties in molecules like those in O2.
AO are associated with the individual atoms, but MO are associated with the whole molecule.
1s 1s
Molecular orbitalsE
*1s
*AO & MO in H2
Atomic orbitals
Anti-bonding orbital
Combination of two 1s AO from each H forms two MO in H2 molecule.
Bonding MO: form between nuclei and are stable Antibonding MO: marked with *; form “behind” nuclei
and are less stable.
1s 1s
1sBonding orbital
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43Energy-Level Diagrams:
44Energy-Level Diagrams:
45Bond Order stability of a covalent bond
Bond Order = ½(# of bonding e- - # of antibonding e-)
Bond Order of 1 = single bond
Bond Order of 2 = double bond
Bond Order of 3 = triple bond
Bond Order of 0 = no bond exists
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46Second Row Diatomic Molecules 1s orbitals combine to form δ1s
and δ*1s 2s orbitals interact to produce δ2s
and δ*2s
δ2s and δ*2s have a greater energy gap
6 electrons must fill Li2 MO
Bond order is 1 δ1s and δ*1s completely filled
therefore contribute little to bonding
Illustrates general rule that core electrons do not participate in bonding
47
Draw the Molecular Orbital for Be2? What is the bond order? Does this compound exist according to the MO diagram?
48*Types of MO
Sigma () MO: form from combinations of: Two 1s or 2s orbitals from different atoms; written
as 1s or 2s. Two 2pz orbitals from different atoms (axial overlap);
written as 2pz.
Pi () MO: form from combinations of: Two 2px or 2py orbitals from different atoms; written
as 2pxor 2py
. Do not appear until B2 molecule
17
49*MO diagrams for “< O2”
Bond order =
½ (# bonding e-
Resulting MO for diatomic molecules with < 16 e- (B2, C2, N2, etc.)
- # antibonding e-)
B.O. (N2) = ½ (10 – 4) =6 / 2 = 3 (triple bond)
N2 has no unpaired electrons which makes it diamagnetic.
