Chemistry 103A; Sections 5, 6, 7, 8; Lecture 30, 3 Nov 00

Recall:

The arrangement of the lone pairs and s-bonded pairs around a nucleus is determined simply by the pairs trying to reduce crowding. That is, the pairs try to get as far away from each other as they can.

We have already seen how two, three, and four pairs arrange themselves. Now we go on the expanded octet arrangements.

(Fortunately, most of the molecules that we work with obey the octet rules. a very very small fraction of molecules have and expanded octet. And most of these are atoms in the second and higher rows bonded to the very electronegative elements, like F and O.)
  Five pairs: the electron pairs are pointing toward the corners of a regular trigonal bipyramid. A five electron pair molecule can have any one of four shapes depending on how many of the pairs are lone pairs.  
PF5, SF4, ClF3, XeF2  
(Use the principle that lone pairs take up more space than s-bonded pairs to determine the shape of the molecule.) Six pairs: the electron pairs are pointing toward the corners of a regular octahedron. (Or, you can think of them as pointing toward the faces of a cube, with the central atom at the center of the cube.) A six electron pair molecule can have any one of five shapes depending on how many of the pairs are lone pairs.  
SF6, BrF5, XeF4


(Recall that lone pairs take up more space than s-bonded pairs.)

 Bond Angles

It takes three nuclei to form a bond angle. (The angle between a s-bonded pair and a lone pair is not a bond angle because there is no second nucleus attached to the lone pair.)

The nominal bond angles produced by the various arrangements of electron pairs are:

linear = 180o

trigonal planer = 120o

tetrahedral = 109.5o (Calculating this is in interesting problem in solid trigonometry. If you have a good background in math, you might want to try to flex your mathematical "muscles" by calculating this one.)

trigonal bipyramidal = (some are 120o and some are 90o)

octahedral = 90o


Bond Angles When Molecule Has Lone Pairs

Lone pairs, of course, do not have bond angles, but their presence in a molecule can influence the bond angle formed by s-bonded pairs.

The principle is that lone pairs take up more space than s-bonded pairs. So the presence of lone pairs may crowd the s-bonded pairs into a smaller space than we would expect and thus decrease the bond angles of the s-bonded pairs.

Examples:

CH4 = 109.5o
In CH4 all the pairs are s-bonded pairs. No one pair takes up more space than any other pair.
NH3 = 107.5o
In NH3 one of the pairs is a lone pair. This lone pair crowds the three s-bonded pairs together slightly and reduces the nominal bond angle of 109.5o to 107.5o. H2O = 105o  
In H2O there are two lone pairs which crowd the two s-bonded pairs even more to reduce the nominal 109.5o ever further to 105o. The order of repulsive interaction of electron pairs is that

lone pair - lone pair > lone pair - s-bond pair > s-bond pair - s-bond pair.

This same effect alters the bond angles in molecules with expanded octets.
 

Molecular Polarity

Now that we know the shapes of molecules we can deduce the polarity of a molecule from the polarity of the individual bonds in the molecule.

The principle is that dipole moments are vector quantities. When we add up the dipole moments of the individual bonds we must use vector addition.

We will do some practice of simple vector addition on the board.

Vectors in line (linear)
CO2, HCN
Trigonal planar vectors (also bent molecules)
SO3, SO2
Tetrahedral vectors (also trigonal pyramidal and bent molecules)
CH4, NH3, H2O, HF
Trigonal bipyramidal and octahedral vectors.
PF5, SF6