Simple bonding models and tools for
predicting shapes: Lewis structures,
VSEPR and hybridization
Angus P. Wilkinson
CHEM 1311A Spring 2004
See “Winter” Chapters 2, 4 and 5
and “Norman” Chapter 4
Structures and bonding models for covalent
compounds of p-block elements
Lewis structures show
– Atom connectivities
– All valence electrons of each atom in a molecule or ion – Formal charges consist with the number of electrons
assigned to each atom
Valence Shell Electron Pair Repulsion
– Foolproof method of predicting molecular structure
based upon the number of σbonding electron pairs and lone pairs; π bonding pairs do notinfluence the gross structure
Hybridization
– Provides equivalent orbitals for all σbonding electron
Structure and bonding, continued
Prediction of composition and structure of binary compounds
– Start with valence shell electron configurations of combining atoms and the notion that a chemical bond consists of a shared electron pair in a bonding orbital formed by overlap of orbitals on adjacent atoms. Overlapping orbitals may be atomic or hybrid orbitals as appropriate
»Formed by combination of two atoms each with a singly occupied orbital »Or by combination of a filled orbital on one atom with an empty orbital
on another
»Utilize valence orbitals as required up to maximum allowed for period
Bonding
There are many different approaches to describing
bonding
– for elements with a large electronegativity difference it is often easiest to view the bonding as arising from
electrostatic interaction between ions formed by the complete transfer of electrons between atoms
» Ionic bonding
– when the electronegativity difference is not so large bonding may be viewed as covalent and involving the sharing of electrons by bonded atoms
Covalent bonding
Modern methods for describing bonding make use of
quantum mechanical methods and describe the
electrons in molecules in terms of molecular orbitals
However, simpler earlier theories such as that due to
Lewis can be useful in some cases and they are a lot
less complex
Lewis theory
Lewis’ theory of bonding was one one the
earliest to have any success
Based on the octet rule
– main group elements like to have eight electrons
when they form compounds (except hydrogen)
Works quite well for second period elements
Simple examples
O
H
H
H
N
H
H
P
Cl
Cl
Cl
Bond order
In many cases Lewis structures can be used
to calculate bond orders that correlate well
with experimentally measured bond
strengths and lengths
Formal charges
A molecule may have more than one plausible
Lewis structure
The best Lewis structure is the one that has the
least charge separation and puts the negative
charge on the most electronegative elements
N N O
N N
O
N N O
-1 +1 0 0 +1 -1 -2 +1 +1
Arrive at formal charges by dividing shared electrons equally amongst the atoms they are shared over
Resonance structures
Molecules and ions with more than one distinct
but equivalent (by rotation or reflection etc.)
Lewis structures can occur. The real structure is
an average of these different resonance forms
Failures of the Lewis model
A number of molecules with odd numbers of
electrons exist (no octet) e.g. NO
An atom may not have enough electrons to
complete its octet without having ridiculous
formal charges e.g. BF
3
A central atom may clearly have more than 8
electrons e.g. SF
6
O
2is paramagnetic!!
VSEPR
Valence Shell Electron Pair Repulsion theory
is a useful tool for predicting the geometrical
structures of main group compounds
Based on the idea that pairs (or groups) of
valence electrons (either bonding or lone
pairs) will try and avoid each other as much as
possible
Groups of electrons and geometry
Electron pair count and geometry
Two pairs on central atom - linear
Three pairs on central atom - trigonal
Four pairs on central atom - tetrahedral
Five pairs on central atom - trigonal
bipyramidal
Six pairs on central atom - octahedral
Electron counting
For molecules with single bonds to the
central atom count all valence electrons of
central atom plus one electron for each
ligand atom
– BCl
33 electrons from boron and 1 from each
chlorine so there are 3 pairs of electrons
When there are double bonds present count
the four electrons associated with the
double bond as a single group
Linear geometry
Molecules with two groups of electrons
Trigonal geometry
Repulsions amongst electron pairs/groups are not all the same. Lone pairs require more space than bonding pairs. One non-bonding
electron requires less space than a bonding pair.
Tetrahedral geometry
Based on a tetrahedral arrangement of
electrons but...
– molecular shape is the arrangement of atoms not electrons
TBP geometry
Lone pairs go equatorial to minimize repulsion
between lone pairs and other groups of electrons
axialequatorial
(b) preferred over (a)
Saw horse Trigonal
bipyramidal
T-shaped
linear
Octahedral geometry
Lone pair avoid each as much as possible
Valence Shell Electron Repulsion Model
SF6 IF5 XeF4 Octahedral square pyramidal square planar 6 bp5 bp, 1lp 4 bp, 2 lp octahedral, 90°
6
PF5
SF4 ICl3 I3
-trigonal bipyramidal see saw T-shape linear 5 bp
4 bp, 1 lp 3 bp, 2 lp 2 bp, 3 lp trigonal
bipyramidal 120° e-e, 90° a-e 5
CH4 NH3
H2O
tetrahedral trigonal pyramidal
bent 4 bp
3 bp, 1 lp 2 bp, 2 lp tetrahedral 109° 28’ 4 BF3 O3 trigonal planar bent 3 bp 2bp, 1lp trigonal planar, 120° 3 BH2 linear 2 bp linear, 180° 2 Examples Molecular shapes Types of
e-pairs
Geometrical arrangement No. e
-pairs
Electron density maps for (CH
3
)
2
TeCl
2
Density map in the CH3-Te-CH3 plane
Density map in Cl-Te-Cl plane (bisecting the C-Te-C angle) Cl
Te
Cl H3C H3C
CH3 Te CH3 Cl Cl Cl Te CH3 H3C
Species that violate VSEPR
Like nearly all truly useful sets of rules
there are exceptions
Transition metal compounds usually do not
follow VSEPR theory
Species that are sterically crowded often do
not obey VSEPR
» XeF6obeys VSEPR (7 pairs)
» TeCl62-does not (7 pairs but is octahedral)
Orbitals and bonding
Moving away from Lewis
structures, we can go to the next level of sophistication and view a chemical bond as involving a pair of electrons associated with an orbital on one atom overlapping with an orbital on another atom
– Here we have two possibilities.
» 1) Both electrons come from one of the atoms (dative bonding or Lewis acid –Lewis base interaction) » 2) One electron comes from each
atom (typical covalent bond)
Cl H
Hybrid orbitals
In a molecule such as CH
4(methane) all
the CH bonds are symmetry equivalent
(indistinguishable)
– This suggests that there are four equivalent
orbitals on the carbon capable of forming
four equivalent C-H bonds
To get four equivalent orbitals on the carbon
we have to
take combinations of the atomic
s and p orbitals to generate hybrid orbitals
sp
3
hybrid orbitals
z x y z x y z x y x y z z x y z x y z x y z x y
sp
2
hybrid orbitals
three orbitals arranged in a trigonal planar fashion with a perpendicular atomic p orbital
z x y z x y x y z z x y z x y z x y z x y z x y z y x
sp hybrid orbitals
two orbitals arranged in a linear fashion with two perpendicular atomic p orbitals
Table of hybrid types
Tetrahedral Geometry
Methane (CH
4)
The 2s and three 2p orbitals of carbon result
Linear Geometry
Beryllium Hydride
promotion
Linear combination of one 2s and one 2p orbitals results in two sp hybrid orbitals:
sp(1)= 1
2(2s+2pz) sp(2)=
1
2(2s−2pz)
Note: The shape and signs of both hybrid orbitals are identical, only the orientation in space is different
Linear Geometry
Ethyne (HCCH)
The two remaining p orbitals contain one electron each
Trigonal Geometry
Boron Trihydride:
The remaining 2p
zorbital is perpendicular to the
the three hybrid orbitals and is not occupied
Trigonal Geometry
Ethene (H
2CCH
2):
The remaining two 2p
xorbitals is perpendicular to the the
Composition and bonding
s px
180(
sp Group 2
s p p p s p p p
BeCl2
Group 13 (≤octet)
sp2
s px py
120(
Composition and bonding
s p p p s p p p
Group 14 (≤octet)
sp3
s px py pz
109(28'
Composition and bonding
s p p p s p p p
ClCl4, SiCl4but also SnCl2and PbCl2
Group 15 (≤octet)
sp3
s px py pz
Composition and bonding
s p p p s p p p
Group 16 (≤octet)
sp3
s px py pz
Composition and bonding
s p p p s p p p
OH2, SH2etc.
Group 17 (≤octet)
sp3
s px py pz
Composition and bonding
s p p p s p p p
Groups 13-17 (3rd row and beyond; 5 bp)
s + px + py + pz + dz2
90(
120(
(sp2 + pd) sp3d
Composition and bonding
s p p p d d d d d
s p p p d d d d d
PCl5AsF5etc.
Groups 13-17 (3rd row and beyond; 4 bp, 1 lp)
s + px + py + pz + dz2
(sp2 + pd) sp3d
Composition and bonding
SF4 etc.
s p p p d d d d d
Groups 13-17 (3rd row and beyond; 3 bp, 2 lp)
s + px + py + pz + dz2
(sp2 + pd) sp3d
Composition and bonding
BrF3etc.
s p p p d d d d d
s p p p d d d d d
Groups 13-17 (3rd row and beyond; 6 bp)
s + px+ py+ pz+ dz2+ dx2-y2 sp3d2
90(
Composition and bonding
SF etc
s p p p d d d d d
Groups 13-17 (3rd row and beyond; 5 bp, 1 lp)
s + px+ py+ pz+ dz2+ dx2-y2 sp3d2
Composition and bonding
BrF5, ICl5etc
s p p p d d d d d
s p p p d d d d d
Groups 13-17 (3rd row and beyond; 4bp, 2 lp)
s + px+ py+ pz+ dz2+ dx2-y2 sp3d2
Composition and bonding
ICl4-etc.
s p p p d d d d d
Rules for assigning oxidation states
The sum of the oxidation states for all atoms in a species must
equal the charge on the species.
The oxidation state of an atom in an elemental form is zero. The oxidation state of a monatomic ion is equal to its charge. Terminal atoms (or groups of atoms) are assigned an oxidation
state consistent with the charge in their monatomic anionic form, i.e., filled valence shell.
The most electronegative element is assigned a negative
oxidation state in compounds of hydrogen.
In compounds such as HO-OH, H2N-NH2, etc. the E-E bond
does not contribute to the oxidation state of either atom. However, in compounds such as SSO32-(S
2O32-, thiosulfate) assignment of a filled valence shell to the terminal sulfur gives it an oxidation state of -2.
Structure and bonding up to now
A chemical bond consists of a shared electron pair in a bonding orbital
formed by overlap of orbitals on the connected atoms
Formed by combination of two atoms each with a singly occupied orbital Or by combination of a filled orbital on one atom with an empty orbital
on another
Composition can be predicted by consideration of the number of singly
occupied orbitals in ground state or low-energy excited state structures of atoms
The geometrical shape of molecules can be predicted on the basis of the
number and arrangement of valence shell electron pairs involved in sigma bonding
The molecular shape is that of the atoms
All of the bonds in a molecule such as CH4are equivalent (identical)
although carbon uses both s and p atomic orbitals for bonding. One model that accounts for this is hybridization