Chapter 10

VSEPR - Lewis structures do not help us predict the shape or geometry of molecules; only what atoms and bonds are involved. To predict shape another system was developed that works for many molecules; Valence Shell Electron Pair Repulsion Theory. The name is derived from the idea that electron pairs around the central atom (the valence electrons) want to be as far away from each other (because of electron repulsion).

Rules: 1. Draw Lewis Structure 2. Count the # of electron pairs around central atom (both bonding & non-bonding) Bonding – electrons that are part of a bond. Non-bonding – electrons that are not in a bond (also called lone pairs) (also, for this method we count multiple bonds as 1 pair)

3. This # determines the basic shape to begin with. 2 e-pairs -> linear - bond angle = 180  Bond Angle – The angle formed between the central atom and 2 terminal atoms. 3 e-pairs -> trigonal planar (a triangle with central atom at center - bond angle = 120  4 e-pairs -> tetrahedral (illustrate) with central atom in the middle - bond angle =  Shape is described this way (i.e. pretending to connect the terminal atoms)

4. Place e-pairs (with atoms where applicable) at each vertex of the geometric shape. 4. Actual shape of molecule is determined by the position of atoms not electrons. Important to understand that the actual shape is determined only by actual atoms. We “see” only atoms, not electrons.

An interesting point is that lone pairs take up more space than bonding pairs. This may result in bonded atoms being pushed closer to each other (since bonded pairs take up less space, they can be closer to each other. An example is H 2 O. The bond angle should be , (based on 4 electron pairs around O, but only 2 atoms). The actual bond angle is about 105 . Let’s do some examples. Note that 5 and 6 e-pair systems exist but we won’t worry about them for now. We will only consider systems with a maximum of 4 electron pairs around the central atom. Let’s do example 10.1, on page 408 (a,b,c and e)

We learned earlier about polar and non-polar bonds. When a molecule is composed of several bonds, the molecule can also be polar or non-polar. The definitive diagnosis is whether or not the molecule has a dipole moment. For our purposes, this simply means that if there is a net positive end and a net negative end for the molecule, it has a dipole moment and is polar.

Many factors contribute to whether a molecule as a whole is polar or non-polar. If the molecule is composed of all non- polar bonds, than obviously the molecule, as a whole, will be non-polar. But even when there are polar bonds in the molecule, it can be non-polar as a whole. This occurs when the shape of the molecule is perfectly symmetrical, and therefore the polarity of the individual bonds cancel out. Some helpful hints will make your job easier:

My rules for non-polar 1.If all bonds are non-polar, molecule is non-polar 2.All hydrocarbons are non-polar A hydrocarbon is any molecule consisting of C and H and nothing else. This is very common as you will see in organic chemistry. 3.If one of the basic shapes, & if every vertex occupied by an identical atom, then non-polar 4.If all else fails, does it look symmetrical?

The Lewis structure application along with the VSEPR model works quite well in predicting structures but to chemists does not really go far enough. They do not really explain why chemical bonds exist or why CH 4 has 4 equal bonds forming a perfect tetrahedron. There are 2 Quantum Mechanical models that help here. We will look at one of them, the Valence Bond Theory.

Basically, in the valence bond model, every atom in a molecule has atomic orbitals just as they would as free atoms. Bonding occurs when an atomic orbital on one atom overlaps with an atomic orbital of another atom. As long as there is significant overlap (technically are the energies with the overlap lower than the separate atomic orbitals) then bonding will occur. It occurs at the point of maximum overlap. See figure 2-4 below

In order to account for molecular geometry it frequently becomes necessary to define some of the atomic orbitals around the central atom differently. It has been discovered through experimental evidence that when an atom, such as C bonds to other atoms, in order to obtain maximum overlapping, the C atom combines s and p orbitals to form new hybrid orbitals. This process is called hybridization. These depend on how many atoms are bonding to the C atom:

4 atoms - sp 3 – It is named because it involves mixing of one s orbital with 3 p orbitals The shape of these new orbitals will be tetrahedral. See Figure below:

The picture above shows how 4 sp 3 bonds to 4 H atoms using the 1s orbitals of the H atoms, forming 4 identical bonds.

3 atoms - sp 2 - trigonal planar and will usually involve one double bond. The second bond of a double bond is called a pi bond and has less overlap that a sigma (first bond) bond and is therefore weaker. It results from the parallel overlap of p orbitals on both atoms:

2 atoms - sp - linear and will usually involve one triple or 2 double bonds. We won’t worry about hybridization involving d orbitals

In valence bond theory, the first bond between 2 atoms is the strongest, because of maximum overlap and are called sigma (  ) bonds. Any additional bonds between the same 2 atoms are weaker because of reduced overlap and are called pi (  ) bonds.