Molecular Geometry and Bonding Theories
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Transcript of Molecular Geometry and Bonding Theories
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- pairs A single, double, or triple bond counts as one domain
Summary of LmABn (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
3Tables 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
Linear
AB2
Linear
180ºBeF2
CO2
A|X
X
A|B
B
3
Three sp2 hybrid orbitals
Trigonal planar
AB3
Trigonal planar
120º
BF3
Cl-C-Cl< 120º
Cl2CO
LAB2
Bent
< 120º
NO21-
A|X
XX
A|B
BB
A|B
:B
5Example: CH4
Molecular shape = tetrahedral
Bond angle = 109.5º
H|
H—C—H|H
109.5º
109.5º
109.5º
109.5º
4
Four sp3
hybrid orbitals
or
Tetrahedral
AB4
Tetrahedral
109.5º
CH4
LAB3
Trigonal pyramidal
< 109.5º
Ex: NH3 = 107º NH3
L2AB2 Bent
<<109.5º
Ex: H2O = 104.5º
H2O
X
A
X
XX
X
A
X
XX
B
A
B
BB
:
A
B
BB
:
A
B
B:
7PCl5
Molecular shape = trigonal bipyramidal
Bond angles
equatorial = 120º
axial = 90º
:Cl: :Cl:\ /
:Cl—P—Cl:|
:Cl::
:::
:: :
120º
120º
90º90º
90º
5
Five sp3d
hybrid orbitals
Trigonal bipyramida
l
AB5
Trigonal bipyramida
l
Equatorial = 120º
Axial = 90º
PCl5
LAB4
Seesaw
Equatorial < 120º
Axial< 90º
SF4
X
X
X
A
X
|
X
B
B
B
A
B
|
B
:
B - A - BB B
5
Five sp3d
hybrid orbitals
Trigonal bipyramidal
L2AB2
T-shaped
Axial<< 90º
ClF3
L3AB2
Linear
Axial = 180º
XeF2
X
X
X
A
X
|
X
B
:
:
A
B
|
B
:
:
:
A
B
|
B
129.3: 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 EN between atoms in bond. Partial (+) and partial (-) on atoms in a polar
bond can be represented as + and -.
H-Cl:
::
H-Cl:
::
+ -
Bond polarity is most often represented by an arrow that points toward the - (most EN atom), showing the shift in e-
density.
13
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).
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.3. Add the vectors, and draw the overall dipole
moment. If none, then = 0.
= Q r
Ex: Draw Lewis structures, bond dipole moments, and overall dipole moments. Also, name the e- domain geometry and the molecular geometry.
CO2 BF3 H2O
CCl4 NH3 PH3
159.4: 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
(kJ/mol)0
-436
0.74 ÅH-H distance
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 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
169.5: Hybrid Orbital Theory Explains the relationship between overlapping
orbitals (valence bond theory) and observed molecular geometries (VSEPR theory).
17“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.)
Be (ground state) → Be (promoted) → Be (sp hybrid)
2p
2s
Energy →
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.
F F
19“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 the s and p e- can’t make the observed 120º angle.
2p
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.
B (ground) → B (promoted) → B (sp2 hybrid)
F F F
20“sp3” hybrid orbitals CH4 (g): observed as tetrahedral
2p
2s
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.
C (ground) → C (promoted) → C (sp3 hybrid)
H H H H
21“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
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.
P (ground) → P (promoted) → P (sp3d hybrid)
Cl Cl Cl Cl Cl
22“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.
23Non-bonding e- pairs Lone pairs occupy hybrid orbitals, too
Ex: H2O (g): observed as bent; but e- domain is tetrahedral
2p
2s
Four sp3 orbitals (2 bonding, 2 non-bonding)
O (ground) → O (sp3 hybrid)
2 non-bonding pairs(lone pairs)
2 bonding pairs
H H
249.6: Multiple Bonds
Draw Lewis structures. For C’s: label hybridization, molecular geometry, and unique bond angles
C2H6
C2H4
C2H2
C6H6
25Sigma and Pi bonds
Sigma () bond: Covalent bond that results from axial overlap of orbitals
between atoms in a molecule Lie directly on internuclear axis “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
26Sigma () bonds in C2H4
Ex: ethene; C-C and C-H -bonds result from axial overlap of H s-orbitals and C sp2-orbitals
27Pi () bonds in C2H4
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
2pC
2s sp2 hybrids bond axially = bonds
p orbital bonds side-by-side = bond
28Sigma () bonds in C2H2
Ex: ethyne (a.k.a. acetylene) C-C and C-H -bonds result from axial overlap of H s-orbitals and C sp-orbital
29Pi () bonds in C2H2
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
sp hybrids bond axially = bonds
2pC
2s
p orbital bonds side-by-side = bonds
30Sigma () bonds in C6H6
Ex: benzene; C-C and C-H -bonds result from axial overlap of H s-orbitals and C sp2-orbitals
32Delocalized 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.
339.7: 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, 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.
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
Molecular orbitalsE
1s
*1s
*AO & MO in H2
Atomic orbitals
Bonding orbital
Anti-bonding orbital
35*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
2px or 2py
.
Do not appear until B2 molecule
36*MO diagrams for “< O2”
Bond order =
½ (# bonding e-
- # antibonding e-)
B.O. (N2) = ½ (10 – 4) =6 / 2 = 3 (triple bond)
N2 has no unpaired electrons which makes it diamagnetic.
Resulting MO for diatomic molecules with < 16 e- (B2, C2, N2, etc.)
N atom N atom
37*MO diagrams for “≥ O2” Resulting MO for
diatomic molecules with ≥ 16 e- (like O2, F2, Ne2, etc.)
O atom O atom
Bond order =
½ (# bonding e-
- # antibonding e-)
B.O. (O2) = ½ (10 – 6) == 2 (double bond)
O2 has unpaired electrons which makes it paramagnetic.
39Magnetism
In an element or compound: Diamagnetism: all e-
paired; no magnetic properties Paramagnetism: at least 1 unpaired e-
* Drawn into exterior magnetic field since spins of atoms become aligned; unlikely to retain alignment when field is removed
Ex: N OSc MnO2
* Ferromagnetism: occurs primarily in Fe, Co, Ni Drawn into exterior magnetic field since spins of atoms
become aligned; very likely to retain alignment when field is removed (i.e., “a permanent magnet”)
Nd2Fe14B is very ferromagnetic