Chemistry 103A; Sections 5, 6, 7, 8; Lecture 29, 1 Nov 00

Recall:

Molecular Shapes

The shape of a molecule is very important in determining the chemical and physical properties of the molecule (compound).

We are using the VSEPR method, which stands for Valence Shell Electron Pair Repulsion to determine the shapes of molecules.

The main idea of VSEPR is that electron pairs are pretty stable and do not like to be crowded. Electron pairs try to get as far away from each other as possible.

We use the Lewis dot formulas to tell us how many electron pairs there are and where the pairs are.

Around any given atom in a molecule we have three "types" of electron pairs:

Lone pairs - these are electron pairs which are not involved directly in bonding.

s -bonded pairs - these are electron pairs which form a single bond between two atoms. s -bonded pairs also are part of a double bond or triple bond.
 

One of the electron pairs in a double bond is a s -bonded pair.
  p -bonded pairs - the second pair of electrons in a double bond forms a p -bond. In a p -bond the electrons density is not between the nuclei, but off of the internuclear axis.

In a triple bond one of the pairs forms a s -bond and the other two pairs of electrons form TWO p -bonds.

 The principal of VSEPR is that lone pairs and s -bonded pairs (including the s -bond parts of double or triple bonds) will try to arrange themselves as far away from each other as possible.

What about resonance hybrids? Use any one of the contributing structures to determine the lone pairs and s -bonded pairs.

(Actually, you can use the resonance hybrid itself because the important lone pairs and s -bonded pairs are the same in the resonance hybrid as in the contributing structures. The electrons that "resonate" are all in p -bonds.) In summary, to determine the shape of a molecule using VSEPR we count only lone pairs and s -bonded pairs.

The rest is simple geometry.

Two pairs will arrange themselves on opposite sides of their nucleus to give a linear arrangement (180o apart).

Three pairs will arrange themselves equally spaced in a plane to give a trigonal planar arrangement (120o apart).

Four pairs will arrange themselves in a tetrahedral arrangement (109.5o apart).

 For an expanded octet we have the following arrangements. Five pairs will arrange themselves to form a trigonal bipyramidal arrangement (90o or 120o apart) Six pairs give an octahedral arrangement (90o apart).

CAUTION: This is a description of what the electron pairs are doing. The arrangement of the electron pairs is not necessarily the shape of the molecule.

The shape of the molecule is determined by where the nuclei are. The s -bonded pairs point to the nucleus of the other atom in the bond.

The lone pairs are not involved in a bond so that there in no second nucleus attached to a lone pair.

 Lone pairs help determine how the electron pairs are arranged in space, but only s -bonded pairs determine the shape of the molecule.

Examples:

Two pairs: linear arrangement of electron pairs.
CO2, C2H2, HCN
Three pairs: the pairs are trigonal planer, the molecule may be trigonal planer or bent.  
H2CO, SO3, SO2
  Four pairs: the electron pairs are pointing toward the corners of a regular tetrahedron, the molecule may be tetrahedral, trigonal pyramidal, or bent.  
CH4, NH3, H2O
  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 that 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.