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Molecular Geometry

This is the second part of a 2 part series on how to determine molecular geometry.  If you haven't read it yet, part 1 focuses on VSEPR theory and how to determine electron geometry.  You can see part 1 by clicking here.

Molecular Geometry

This is the easy part!  We got all the grunt work out of the way learning electron geometry, now we just need to apply our knowledge.  You just need to know a couple of things to do this.

1 - In all those pictures, we referred to Y as a magnet.  It is actually either a bond between atoms, or a lone pair of electrons.

2 - When we name molecular geometries, we ignore lone pairs of electrons, even though they are still there and still push away other electrons like the little magnets that they are.

3 - Lone pairs of electrons are stronger magnets than bonds are, so they will push harder.

4 - When every Y stands for a bond, the molecular geometry is the same as the electron geometry.

Like before we will start with the linear electron geometry.  Good news, it does not matter if Y is a bond or a lone pair - they still both from a linear molecular geometry.  Even though lone pairs push harder than bonds, its impossible to get farther away then 180°.  All of these pictures were created using ACD Chemsketch and Chem3D - free chemistry drawing programs.

linear electron geometry       linear molecular geometry

                                             Linear

Next up, we have the trigonal planar electron geometry.  If all of our Y's are bonds, the molecular geometry will again be trigonal planar.  However in this case, we can replace one Y with a lone pair of electrons to get a "bent" molecular geometry.  Just imagine in your head removing one of the bonds and replacing it with the lone pair.  Like rule #3 says, lone pairs push harder than bonds do, so while the angle between bonds is 120° for a trigonal planar molecule, it is actually <120° for a bent molecule, all because that lone pair is pushing the bonds farther away.

trigonal planar electron geometry  

trigonal planar molecular geometry   bent molecular geometry   

                           Trigonal Planar                        Bent

It starts to get fun here with the tetrahedral electron geometry.  Again, if all the Y's are bonds, the molecular geometry is just tetrahedral again with angles equal to 109.5°.  Now if we replace just one of the bonds with a lone pair, we get what is called trigonal pyramidal - it looks like a pyramid with 3 sides - and because the lone pair is pushing harder than the bonds, the angle is <109.5°.  Finally, with a tetrahedral, we can replace 2 of the Y's with lone pairs (i.e. water - 2 bonds and 2 lone pairs) and we get a familiar looking "bent" shape, only the angle between the bonds is now <109.5°.

tetrahedral electron geometry  tetrahedral molecular geometry

                                                                  Tetrahedral

trigonal pyramidal molecular geometrybent molecular geometry

                      Trigonal Pyramidal                        Bent

Now we have to look at the trigonal bipyramidal electron geometry.  Remember how this was a combination of linear and trigonal planar geometries?  This becomes important when we start adding lone pairs.  Since the lone pairs push harder than bonds do, they want to be farther away from everything.  That means that when we add them, they MUST be added to the horizontal plane (trigonal planar geometry).  This puts the lone pairs as far away as possible (120° angle is farther away than a 90° angle). 

So we begin, as always, when all Y's are bonds, the molecular geometry is the same as always - trigonal bipyramidal - with angles equal to 120° and 90°.  When we replace one bond with a lone pair (in the horizontal plane) we get what looks like a see saw if we turn it sideways so we call that a seesaw molecular geometry with angles <120° and <90°.  Replacing two bonds (in the horizontal plane) with lone pairs of electrons gives us what looks like a T if we turn it sideways, so naturally we call that T-shaped molecular geometry, with angles <90°.  Finally, we can replace all three of the bonds in the horizontal plane with lone pairs, and we're just left with our linear geometry (vertical plane) so naturally we call that a linear molecular geometry with an angle of 180°.

trigonal bipyramidal electron geometry

trigonal bipyramidal molecular geometryseesaw molecular geometry

                    Trigonal Bipyramidal                       Seesaw

t-shaped molecular geometry          linear molecular geometry

                                      T-shaped                      Linear

Last, but not least, we have the octahedral electron geometry.  As always, if all Y's are bonds, than the molecular geometry will be octahedral with angles of 90°.  When we replace one bond with a lone pair of electrons, we get another pyramid shape, but this time with 4 sides.  We call this molecular geometry square pyramidal with angles <90°.  Finally, if we replace two bonds with lone pairs of electrons, the lone pairs must be added on opposite sides because of their more powerful magnetic force - they push harder.  This gives us a molecular geometry where we have 4 bonds in a single plane, so we call it square planar where the angles all equal 90°.

octahedral electron geometryoctahedral molecular geometry

                                                                       Octahedral

square pyramidal molecular geometrysquare planar electron geometry

                Square Pyramidal                        Square Planar

That's it!  You are done.  At this point, you should be able to tell the difference between electron geometry and molecular geometry, as well as how they are related.

Did I miss anything in my explanation?  Still not understand something?  Let me know

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