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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

(2)

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

(3)

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

(4)

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

(5)

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

(6)

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

2

is 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

(7)

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

(8)

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

3

3 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

(9)

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

(10)

TBP geometry

‹

Lone pairs go equatorial to minimize repulsion

between lone pairs and other groups of electrons

axial

equatorial

(b) preferred over (a)

Saw horse Trigonal

bipyramidal

T-shaped

linear

Octahedral geometry

‹

Lone pair avoid each as much as possible

(11)

Valence Shell Electron Repulsion Model

SF6 IF5 XeF4 Octahedral square pyramidal square planar 6 bp

5 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

(12)

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

(13)

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

(14)

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

(15)

Table of hybrid types

Tetrahedral Geometry

‹

Methane (CH

4

)

‹

The 2s and three 2p orbitals of carbon result

(16)

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

(17)

Trigonal Geometry

‹

Boron Trihydride:

‹

The remaining 2p

z

orbital is perpendicular to the

the three hybrid orbitals and is not occupied

Trigonal Geometry

‹

Ethene (H

2

CCH

2

):

‹

The remaining two 2p

x

orbitals is perpendicular to the the

(18)

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

(19)

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

(20)

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

(21)

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

(22)

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

(23)

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

(24)

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

Figure

Table of hybrid types

References

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