$Molecular Geometry
Molecular geometry as determined by the Valence Shell Electron Pair Repulsion
Model (VSEPR model)
Molecular Geometry = three-dimensional arrangement of atoms in a
molecule/ion, including bond angles, (bond angles in degrees)
Valence electrons = outer electrons, bonding electrons and lone pair electrons
VSEPR model/theory – molecular geometry is that which minimizes repulsion of
valence electrons pairs by having electron pairs as far away from each other as
geometrically possible. (Gillespie/Nyholm theory)
a. Double/triple bonds between 2 atoms treated as one location
b. VSEPR structure for individual resonance structures
A. Central atom has only bonding pairs/ no lone pairs on central atom: AB2 – AB6
(AB7 has several structures)
Electron pairs
arrangement
180
2
3
o
A
BF3
B
120o
120o
A
FBF = 120o
A
B
B
Trigonal planar
120o
4
Molecular Geometry Examples
B A B
BeCl2, HgCl2
linear
ClBeCl = 180o
CH4, NH4+, BF4
HCH = 109.5o
B
109.5o
A
A
B
B
B
tetrahedral
B
o
90
5
A
120o
B
B
A
B
Equatorial
B
B
B
Cis
A
B
axail
SF6, PF6
FSF = 90o
B
A
FPF = 90o
B
Trigonal bipyramid
6
PF5
FPF = 120o
B
octahedral
FSF = 180o
trans
Relative strength of electron-pair repulsion:
Lone pair vs lone pair > lone pair vs bonding pair > bonding pair vs bonding pair.
(Lone pair “need more space”; bonding pairs contained between atoms)
ABxEy, A = central atom, B = bonded atoms, E = lone pairs on A:
Overall geometry of electron pairs = x + y positions
Geometry of molecule/ion = X positions only; only atoms are measurable.
B. Central atom with Bonding Pairs (B) and Lone pairs (E):
Class, ABxEy
Number
of elec.
pairs
Bonding Lone
pairs
pairs
AB2E
3
2
1
SO2, O2
Overall
geometry
Trigonal
planar
BAB
AB3E
Molecular Geometry
A
B
B
Bent molecule
120o
Tetrahedral
4
3
1
BAB
A
109.5o
B
B
NH3
B
Trigonal pyramid
AB2E2
Tetrahedral
4
2
2
BAB
B
o
109.5
H2O
A
B
Bent molecule
AB4E
SF4, IF4+,
XeO2F2
5
4
1
Trigonal
bipyramid
Bond angles
120o and 180o
B
B
A
B
B
See-saw molecule
AB3E2
5
3
2
ClF3
Trigonal
bipyramid
BAB
90o
B
A
B
B
T-shaped molecule
AB2E3
5
XeF2, I3
2
3
Trigonal
bipyramid
BAB
180o
B
A
B
Linear
AB5E
Octahedral
6
5
1
BAB
o
90
B
B
B
BrF5, XeOF4
A
B
B
Square pyramid
AB4E2
XeF4, ICl4
Octahedral
6
4
2
BAB
90o
B
B
A
B
B
Square planar
Examples:
i
H
N
H
H ii
H
NH3
N
H
H
AB3E, Overall = tetrahedral; Molecular = trigonal pyramid
Repulsion i > ii, so the HNH angle shorter than that normally expected for a
tetrahedral structure (HNH = 107.3o)
H
H2O
O
H
AB2E2, overall geometry = tetrahedral; molecular
i
O
H
geometry = bent molecule
Electron repulsion, I > ii > iii
Actual bond angle HOH is 104.5o
iii
ii
H
F
F
F
S
F
F
SF4
AB4E, 5 pairs  trigonal bipyramid
bond angle, expected is 120o, but actual one is 116
F
S
F
F
FSF
S
O
SO2
molecule.
O
or
O
S
O
both AB2E. Trigonal planar/bent
S
O
O Actual bond angle OSO is 119.5o
Review Procedure
1. Write Lewis structure
2. Count bonding pairs (B) and lone pairs (E) around central atom (A), treating
double/triple bonds as 1 bonding pair, to obtain ABxEy formula.
3. Predict overall geometry of electron pairs from x + y.
4. Predict geometry of molecule by B positions in ABxEy, remembering that one
pair repulsion stronger than bonding pair repulsion.
ICl4 5 x 7 + 1 = 36 electrons/18 pairs
AB4E2 overall = octahedral
Molecular = square planar
ClICl = 90o
Cl
Cl
I
Cl
Cl
Polar (covalent) Bonds/ Dipole Moments
Polar covalent bonds result when electron-pairs are shared between atoms of
different electeronegativities (A), and have a partial separation of charge
(because of this unequal sharing)
+
H

F
dipole molecule align in an electrical field
Dipole moment,  = Q x r
|charge| in coulombs, distance in m
Debye unit (D), 1D = 3.33 x 10-30 C.m
+
+
+
- +
- +



{electron-positron 100 pm separated   = 4.80 D}
Covalent bond between atoms of same element is non polar, because of equallyshared electron pairs: H-H, F-F etc.
Homo-nuclear diatomic molecules,  = 0
Hetero-nuclear diatomic molecules   0 (have dipole moment)
Polyatomic molecules have dipole moment,   0, if:
1. has polar covalent bonds;
2. Individual bond moments do NOT cancel in vector addition

+ +

O
C
O
CO2, AB2, linear molecule,
=0,  = 0
,
Non-polar molecule with polar bonds
+

+
+


S
O
O

SO2, AB2E bent molecule,
polar molecule
,   0, has dipole moment, is



N
+
NH3, AB3E, trigonal pyramid,
NF3,
H
H
H
+
+
+
N

+

F 
F
F



,  0, has dipole moment.
ClHC=CHCl, 1, 2, dichloroehtylene (planar molecule)
,   0

H
C
H
C
Cl
Cl
,  0
cis-isomer
H
Cl
C
C
H
Cl
trans-isomer
 = 0, no dipole moment
When all the B atoms the same, the following ABxEy formulas
 = 0 for AB2, AB3, AB4, AB5, AB6, AB2E3, AB4E2
  0 for AB2E, AB3E, AB2E2, AB4E, AB3E2, AB5E
BF3(AB3),  = 0; NF3(AB3E)   0
Care should be taken when considering different B atoms; can have   0.
BF3, BBr3  = 0

F

B
But, BFBr2,   0, Br
Br
Valence Bond Theory
Overlap of atomic orbitals illustrates covalent bonds; H-H
+
H 1s-1s
-bond
covalent bond
- bonds – end-on overlap of atomic orbitals along bond axis:

+
+
F-F, [He]2s2 2px2 2py2 2pz1
same sign
+
H-F
, 2pz-2pz, overlap of

+
2pz
1s

Multiple bonds
1  – bond
2  - bonds
1s-2pz
N
N
[He] 2s2 2px1 2py1 2pz1
- bond – parallel or side-by-side overlap of orbitals producing an electron cloud
above and below bonding axis but a nodal plane through the axis.
+
+
+



one -bond
Single bond =  – bond, H-H, F-F, H-F etc
Double bond = 1 -bond + 1 -bond, O=C=O etc
N N
Triple bond = 1 -bond + 2 -bonds,
-bonds usually slightly weaker than -bonds
Hybridization of Atomic Orbitals
Simple atomic orbitals often suggest wrong number and geometry of bonds
e.g C [He]2s2 2px1 2py1 – suggests 2 covalent bonds at 90o angle- but CH4, in
general CX4, tetrahedral – is observed
Hybridization is a possible explanation
1. Promotion (excitation) of valence electron:
[He]2s2 2p2  [He] 2s1 2px1 2py1 2pz1 – an allowed excited state that would give 4
covalent bonds, but wrong angles and unequal bonding
2. Hybridization (mixing or linear combination) of valence atomic orbitals
[He] 2s1 2px1 2py1 2pz1  [He] four (sp3 hybrid orbitals)4
Shape of each sp3 hybrid orbital:
4 equivalent sp3 hybrid orbitals directed to vertices of tetrahedron
CH4
H
C
H
H
H
Four  (Csp3 – H1s) bonds
Hybridization of 4 atomic orbitals (S, px, py, pz) produces 4 hybrid orbitals (sp3)
having correct geometry (tetrahedral) for covalent bonds on central atom with
equivalent bonding.
n atomic orbitals  n hybrid orbitals
NH3 – use 4 sp3 hybrid orbitals
lone pair occupies
non-bonding sp3
NH
H
H
H
N – [He] 4 (sp3)5
bond angle – 107.3o
H2O – use 4 sp3 hybrid orbitals: O – [He] 4 (sp3)6 – two covalent bonds + 2 lone
pairs of electrons
lone pair occupies
2 non-bonding sp3
O
H
H
Sp hybridization
BCl2(g) – linear molecule
Be – [He] 2s2 – promotion  [He] 2s1 2pz1
sp linear combination
Two sp hybrid orbitals – orbitals at 180o angle
Sp2 hybridization
BF3 – trigonal planar
B – [He] 2s2 2px1  [He] 2s1 2px1 2py1  [He] three sp2 hybrid orbitals
Three  – bonds (B sp2 – F2pz)
Three sp2 hybrid orbitals directd to points of triangle (120o angles)
Sp3d hybridization
PF5 trigonal bipyramid
P – [Ne] 3s2 3p3  [Ne] 3s1 3px1 3py1 3pz1 3d1 – [Ne] five (sp3d)5
Five sp3d hybrid orbitals directed to vertices of a trigonal bipyramid ; 5 s bonds
(Psp3d – F2pz)
Expanded octet/expanded valence shell by atoms having available d orbitals (3 rd,
4th period): P, As, Sb, (NOT N), Cl, Br, I (not F)
Sp3d2 hybridization – SF6 octahedral
S: [Ne] 3s2 3p4  [Ne] 3s1 3px1 3py1 3pz1 3x2-y2d1 3dz21
Six  covalent bonds (Ssp3d2 – F2pz)
Summary of Hybridization schemes:
Number of
Valence electrons
2
3
4
5
6
Overall geometry
valence electron pair
Hybridization
scheme
lnear
trigonal planar
tetrahedral
trigonal bipyramid
octahedral
sp
sp2
sp3
sp3d
sp3d2
Hybridization procedure
Draw Lewis structures /determine geometry by ABxEy class
Match geometry to hybridization scheme
Hybridize atomic orbitals on central atom in molecules, never isolated atoms
N atomic orbitals used – n hybrid orbitals formed
Energy needed for promotion recovered in bond formation
Covalent bonds explained as overlap of orbitals