0:02 everyone so up until now we've been
0:03 working with lewis structures
0:05 and the problem with lewis structures is
0:07 that we're not really told
0:10 what a bond actually is so we're going
0:11 to work today
0:13 on valence bond theory so valence bond
0:14 theory is built
0:16 as the first theory from quantum
0:18 mechanics that actually explains the
0:21 origins of where our bonds come from
0:23 so these are our two learning outcomes
0:24 you should be able to describe how
0:27 covalent bonds are formed in terms of
0:29 atomic orbital overlap and you should be
0:30 able to define
0:34 and draw examples of sigma and pi bonds
0:37 so if we start to take two hydrogen atoms
0:38 atoms
0:40 and we start them really far apart and
0:43 we move them closer together
0:45 so we start moving them together they
0:46 will actually
0:49 start to form an energy drop so if we measure
0:50 measure
0:54 energy as we go over distance
0:57 and then we start to look at this from
0:59 down here we bring them closer and
1:00 closer together we'll see that the energy
1:01 energy
1:03 actually starts to fall and when it
1:05 actually falls it actually hits a
1:08 well in here where we start to get some overlap
1:09 overlap
1:12 of our electron density from the two
1:14 hydrogen atoms
1:15 and if you push them closer together
1:17 they start to repel each other the
1:18 nuclei repel
1:20 each other so as you go here energy goes
1:21 in infinite
1:24 but we're interested in this energy well
1:25 right here
1:27 where we actually start to form an
1:28 overlap of
1:32 atomic orbitals so one of the orbitals
1:34 that we can kind of think about
1:37 in here in terms of the uh overlap are
1:38 sigma box
1:40 so this is when we have our two hydrogen atoms
1:42 atoms
1:44 and when we have these two hydrogen atoms
1:45 atoms
1:48 we get this overlap that is along
1:55 so we normally draw our bond as a line
1:57 in between the two nuclei but in the
2:00 sigma bond we end up getting lots of
2:01 overlap with this bond
2:03 so we end up getting something that
2:06 looks like this
2:08 where we get lots of this electron
2:09 density in the middle
2:12 and we will normally draw this as just a
2:16 line on here so that is an example of an
2:19 s to s kind of overlap you can also have
2:20 this with an
2:23 s and p orbital where we draw the line here
2:24 here
2:26 uh for the bond and we see overlap of
2:27 the p
2:29 lobe and the s orbital and we can also
2:30 have p
2:34 to p overlap in our orbitals
2:36 so if we draw our bond on here the two
2:39 lobes actually overlap with each other
2:42 to form lots of electron density along
2:43 the bond itself
2:47 then we come along to a pi bond so the pi
2:47 pi
2:50 bond the electron density is
3:04 so if we have our p orbitals we can
3:07 actually look at them
3:09 so if we have our two atoms on here we
3:10 have our two p
3:13 orbitals these dumbbell shapes
3:16 they can actually start to form overlaps
3:19 with each other so you end up getting
3:21 something that looks like this we draw our
3:22 our
3:24 our bond is typically drawn in here
3:26 you'll end up getting lots of electron densities
3:27 densities
3:32 above and below the bonds themselves so
3:34 you have this nice little bond
3:36 that forms and typically you'll see a
3:37 picture like
3:40 this in your in your books and you'll
3:43 see lots of overlap with these p
3:46 orbitals to form this cloud of electrons
3:50 above and below the bond itself
3:53 so the bond is in here so what can we do
3:55 with this we can actually start to count
3:58 sigma and pi bonds in our different
4:01 lewis structures
4:03 so if we start off with just a single
4:05 bond and this hydrogen chloride
4:09 hcl atom on here and just a single bond
4:12 that is always due to a sigma bond
4:14 if we look at oxygen we draw the lewis
4:17 structure we have a double bond in here
4:20 the double bond in here is actually a sigma
4:21 sigma
4:25 and a pi bond so you have one sigma
4:28 and one pi bond
4:31 on here and then if we look at nitrogen
4:32 which we draw the lewis structure we have
4:33 have
4:38 a triple bond one of these is a sigma
4:40 and then you have two pi bonds so each
4:42 of these lines that we're drawing in our
4:44 lewis structures
4:47 are a sigma and two pi bonds
4:50 so this theory actually tells us what
4:51 these these bonds are
4:53 actually about and we can start to
4:56 sketch them out
4:59 and what i also like about valence bond
5:01 theory is we can start to think about dipoles
5:01 dipoles
5:05 and you know that you know how we share
5:08 our electrons
5:10 so if we start off with something like
5:13 hcl we already know that the chlorine
5:15 side is going to be slightly negative
5:17 the proton side is going to be slightly
5:19 positive and valence bond theory will actually
5:20 actually
5:22 help us to draw that most of the
5:24 electron density
5:26 is going to be around the chlorine so
5:27 you have
5:29 lots of electrons around the chlorine
5:31 which gives it a slightly
5:33 negative charge mostly electron density moves
5:34 moves
5:36 away from the hydrogen so you get a
5:38 slightly positive
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