topic 5 - VSEPR theory

lewis theory

• a cov bond is formed when two neighbouring atoms share an electron pair

• octet rule : atoms share electron pairs until they have acquired an octet of valence electrons - except doublet in H

• only valence electrons are important in making cov bonds

• assumes localised bonding electrons - electrons stay assigned to a given atom in a cov bond

limitations of Lewis theory

• lewis diagrams don’t directly allow the prediction of molecular geometry

• exceptions to octet rule: electron difficient molecules eg. BeH2, BF3 ad hypervalent molecules w expanded octet eg. SF6 and PCl5

What is VSEPR theory?

• valence shell electron pair repulsion theory

• used to predict molecular geometry and deviations from ideal geometry eg. in bond lengths and angles

• works with both electron deficient and hypervalent molecules

primary assumption:

• regions of enhanced electron density stay as far away from each other as possible

• regions= bonding pairs, lone pairs, single unpaired electrons (radicals)

what are the 5 rules of VSEPR theory?

1. stereochemsitry is determined by valence shell electron pairs arranging themselves to minimise repulsions

2. relative strength of main interactions- lp-lp > lp-bp > bp-bp

3. geometry in determined only from the number of sigma pairs and lone pairs (pi bonds are not counted)

4. bonding pair sizes: A(triple)B > A=B > A-B

5. the size of a bonding pair decreases with increasing electronegativity of the ligand

simplified steps to using VSEPR (1-6)

1. central- least electroneg where possible

2. no. of electrons- all in outer s+p orbitals

3. add donated- e- donated to bonding by ligands

4. charge- add for negative, subtract for positive

5. pairs- divide total e- by two

6. distribute ligands- consider pair repulsions + minimise

predicted geometry of electron pair distribution based on 7 valence electron pairs

• pentagonal bipypamidal

• sp3d3 hybridisation of central atom

predicting distortions from ideal geometry- PCl5 - not all the P-Cl bonds are the same length, why?

• shape - trigonal bipyramidal

• P-Cl_ax > P-Cl_eq

• extra repulsion from having 3 atoms at 90 degrees rather than 2 from the equatorial chlorines therefore equatorial bond length is slightly shorter

what are the relative sizes of bonding and lone pairs?

• a bonding pair is attracted to two nuclei

• a lone pair is attracted to only one

• hence the lone pair is effectively more repulsive as it has a greater density closer to the nucleus

• more effective at pushing other atoms away

implications of rule 2: example of SF4

• electron pair distribution shape - trigonal bi pyramidal - has one lone pair so actually sawhorse

• most repulsive interactions occur between pairs at 90 degrees

• so having the lone pair equatorial is more favourable than axial

• all bonds are pushed away from long pair and closer to each other

• by default lone pairs should be placed on the equatorial sites to reduce repulsion

sawhorse shape

• 5 total electron pairs

• 1 lone pair in an equatorial site

T-shaped molecular shape

• 5 total electron pairs

• 2 lone pairs in equatorial sites

linear? (5 pairs)

• 5 total electron pairs

• 3 lone pairs in equatorial sites

square planar (6 pairs)

• 6 total electron pairs

• 2 lone pairs (in axial positions)

how to determine how many electrons ligands donate

• consider how many ligand electrons are involved in bonding to central atom

• lone pairs in single atoms are not counted eg. N and P are 3 e- donors and O, S and Se are 2 electron donors

• do not break bond on multiple atoms ligands eg. CH2 donates 2 electrons form the Cp orbitals

implications of rule three: example NOCl

• 1st bond = sigma

• 2nd bond = pi

• 4 electron pairs : 2 sigma, 1 pi and one lone pair

• electron pair distribution shape = trigonal planar

• molecular shape = bent/non-linear

• distortions likely due to presence of lone pair

• pi bonds are slightly more repulsive than sigma bonds but not more than lone pairs bc electron density isn’t as close as in lone pairs - above + below internuclear axis

• O-N-Cl = <120 degrees

implications of rule 4 : example POCl3

• electron density increases with increasingly multiple bonds

• 4 sigma bonds, 1 pi bond, no lps

• shape of electron pair distribution + molecular shape = tetrahedral

• deviations from ideal geometry likely due to presence of pi bond

• Cl-P-Cl <109.5 degrees

• O-P-Cl > 109.5 degrees

• (not relevant here but lp have less effect on pi bonds than sigma bonds)

do long pairs have a greater effect on sigma or pi bonds?

• sigma

implications of rule 5

• the greater the electronegativity the lower the repulsive nature

• electron density in O-F is closer to F due to high electroneg so F are more able to be pushed together - bonds are less repulsive

• result H-O-H (H2O)= 104.5 vs F-O-F (F2O) = 103.2

• also NH3 vs NFs

• F-N-F bond angle is smaller as N-F bonds can be pushed closer together as there is weak electron density in the bond the closer to N you get

effect of radical species

• a radical species are those w a single unpaired valence electron - very reactive

• in VSEPR a radical e- is treated as a small lone pair - a very weak species

• lp > A(triple)B > A=B > A-B > e-

• larger angle in NO2 compared NO2 - shows weak repulsive nature of radical - pi bonds are pushing back against it increasing the angle past 120 degrees

• radicals still contribute to shape so NO2 has a trigonal planar shape