Sunday, October 09, 2005

Lecture 002

Instructor: Jean-Claude Bradley, Drexel University
View Lecture


1
Chem. 241 - Lecture 2
1 DISCLAIMER: This text is being provided in a
rough-draft fashion. Communication Access
2 Realtime Translation (CART) is provided in order
to facilitate communication accessibility and may
3 not be a totally verbatim record of the
proceedings.
4
* * *
5
6 Lecture 2
7 Okay. So we're going to start with
8 electronic configuration, so this will be a bit of
9 a review for you, but you have to have this
10 material fresh in your mind to actually see how we
11 will implement it. If you remember all this, just
12 bear with it a little bit.
13 You're going to be applying two rules
14 that you should be familiar with the -- the Pauli
15 solution principle -- and, by the way, the
16 software I'm using to write it on, I'll make that
17 available to you as a PDF, so you don't have to
18 worry about copying everything down because it
19 will be available in an easily printable format.
20 So the Pauli solution principle has to
21 do with the fact that only two electrons can
22 occupy one orbital.
23 And the other rule we're going to need
24 to be able to use to deal with electronic
25 configuration of the first ten elements is the
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2
Chem. 241 - Lectures 2 and
1 Hunds rule, where electrons like to be unpaired if
2 possible.
3 Okay. How many of you remember these
4 two rules? Okay, good. So we're going to apply
5 these. And we're basically going to, like I said,
6 fill up the first ten elements. So we're going
7 from hydrogen to neon. And these little lines
8 that I'm drawing are the orbitals, and those will
9 be the orbitals where we have to apply the Pauli
10 solution principle and the Hunds rule.
11 So as we go from the top to the bottom,
12 we're adding one electron per each element. So
13 hydrogen has one electron. And let me use a
14 different color here. We can show that electron
15 with a half arrow either pointing up or pointing
16 down. The direction of the arrow has to do with
17 the skin of the electron, which we don't have to
18 get into; it's just that you need to know that
19 this can either be up or it can be done.
20 So when we move the helium, we have two
21 electrons. Apply the two rules, put the helium.
22 So following the two rules, we put the first
23 electron there and the second electron. It
24 doesn't have another orbital of the same energy,
25 so the other electron can't remain unpaired so
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3
Chem. 241 - Lectures 2 and
1 we're going to be forced to pair it. And the way
2 we demonstrate that is by drawing a down arrow
3 next to the up arrow. So now we have two
4 electrons in the orbital, and it's filled up so we
5 can't put any more in there.
6 What is the name of this first orbital?
7 1S. So obviously, I'm going to have to draw more
8 lines here. I'm going to draw all the lines for
9 all the elements, and I'm leaving them blank until
10 I fill them up.
11 So continuing on with helium, it has
12 three electrons total. So the first one goes in
13 the first orbital. And the second one goes in the
14 same orbital, different direction. And then the
15 third one we put in the next orbital. So the
16 first orbital is 1 and the second orbital is 2S.
17 So now we continue with beryllium.
18 It has four electrons. Now I just filled up
19 the 2S orbital, so when I go to boron, that
20 has five electrons and I now have to use
21 these other three orbitals that are of the
22 same energy, and those would be the 2P
23 orbitals.
24 So I have 2PX, 2PY, and 2PZ. And
25 there's no particular reason for that order;
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4
Chem. 241 - Lectures 2 and
1 it's just distinguished between them.
2 They're absolutely the same energy. So that
3 like electron, then I have to put it in the
4 2PX orbital like that
5 Okay? So now we have one more
6 electron -- one, two, three, four, five --
7 and now we have three orbitals that are the
8 same energy. So the electrons will be
9 unpaired if possible according to the Hunds
10 rule, so we will put them as unpaired. So
11 I'm going to have one in 2PX and one in 2PY,
12 for example.
13 With nitrogen, I have one, two,
14 three, four, five, six, seven. And, again,
15 I'm putting another unpaired electron
16 because it's the same energy. All the 2P
17 orbitals are the same energy.
18 Oxygen -- one, two, three, four,
19 five, six, seven. And now the eighth one,
20 now I'm forced to pair it because I've run
21 out of orbitals of the same energy that are
22 vacant. So I'll put in the 2PX, for
23 example.
24 With fluorine, we fill this up
25 where I'm going to have to pair the electron
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5
Chem. 241 - Lectures 2 and
1 and the 2PY orbitals.
2 And, finally, neon has all of the
3 two P orbitals completely filled up with
4 paired electrons, okay?
5 So this is something I'm taking a
6 bit of time to go through because we will
7 need this when we look at molecular orbitals
8 that is based on this theory. So for now,
9 this should be a refresher.
10 The other reason I go through this
11 is because I want to point out a concept
12 which is valence electrons. When we look at
13 making molecules from collections of atoms,
14 those atoms are connected together with
15 bonds. And it turns out that not all of the
16 electrons in the atoms participate in that
17 bond formation; only the valence bonds
18 participate. So I'm going to put a square
19 around the valence electrons.
20 So, for example, for hydrogen and
21 helium, well, those are all the valence
22 electrons; that's all it has. But when I go
23 from lithium to neon, the 1S orbital is not
24 a part of the valence electrons.
25 So I'm going to put a square
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6
Chem. 241 - Lectures 2 and
1 around just the two orbitals, the 2S and 2P
2 orbitals.
3 So those are called valence
4 electrons. And so if I ask, you know, how
5 many electrons does carbon have? what I'm
6 really interested in is the four electrons
7 that are in the valence.
8 So generally when we talk about
9 how many electrons each atom has, we're
10 referring to valence electrons as we count
11 them
12 Any questions on this?
13 (No questions.)
14 Okay, so the next thing I wanted
15 to with these valence electrons is rewrite a
16 little useful periodic table for you with
17 the valence. The periodic table. So these
18 are the elements that we're going to be most
19 concerned with in this class: Hydrogen,
20 boron, carbon, nitrogen, oxygen, fluorine
21 phosphorous, sulfur, chlorine, bromine, and
22 iodine. We can put even lithium and sodium
23 on the left.
24 The reason this is a valence
25 periodic table is that we're going to divide
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7
Chem. 241 - Lectures 2 and
1 this up into columns and only worry about
2 the number of valence electrons. So
3 hydrogen has one valence electron; boron has
4 three; carbon, four; nitrogen and
5 phosphorous have five; oxygen and sulfur
6 have six. And the halogens -- fluorine
7 chlorine bromine and iodine have -- have
8 seven.
9 Okay? So those elements are going
10 to behave in pretty similar ways. For the
11 two elements have the same number of valence
12 electrons, there will be differences in how
13 they react, but there will be a lot of
14 similarities as well as to how they format
15 bonds with other atoms.
16 So this is something that, you
17 know, you should keep handy when you're
18 doing problems because you're going to need
19 to use it constantly.
20 Okay. So keeping this in mind,
21 let's now talk about actually making
22 molecules from atoms. Well, let's consider
23 the simplest thing we can consider, which is
24 which is two hydrogen atoms coming together
25 to make hydrogen gas, H2.
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8
Chem. 241 - Lectures 2 and
1 So hydrogen has one electron, and
2 I'm getting the Lewis structures here, so
3 this should look pretty familiar. I put
4 hydrogen and then I put a dot, or sometimes
5 we use an X. But the point is, you show
6 that there's only one electron.
7 If I want to show two hydrogen
8 atoms coming together, I can show one
9 electron and the other hydrogen. And then
10 when they come together, I can draw the two
11 electrons on top of each other between the
12 two hydrogens.
13 And what I've done here is, I've
14 created a bond, but there are two kinds of
15 bonds that we're interested in. This is a
16 covalent bond because the two atoms are
17 sharing the electrons. Neither hydrogen has
18 those electrons entirely to itself; it has
19 to share both of them, so that's why it's
20 called a covalent bond.
21 So another way that we can draw
22 the covalent bond is to draw a line. So
23 that's how you join two hydrogens, with a
24 line. Every line you see corresponds two to
25 two electrons. So we will be doing that
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9
Chem. 241 - Lectures 2 and
1 most of the time because it's more
2 convenient than drawing a whole bunch of
3 dots. So we need to draw dots when we don't
4 have bonds; but otherwise, we can just draw
5 a line like this.
6 Let's consider another case
7 lithium fluoride. Lithium has one valence
8 electron. Fluorine has seven valence
9 electrons. But in this case, instead of
10 sharing the electrons, the lithium actually
11 completely gives up its electron, and the
12 fluorine completely takes it, so that's
13 different.
14 If the lithium gives up its
15 valence electron, it now has lost a negative
16 charge, and so it's going to be plus,
17 specifically plus 1. And the fluorine now
18 has eight electrons around it. Fluorine
19 normally has seven electrons, but it has an
20 extra one and so it has an extra negative
21 charge. That's how we know that it's
22 negative.
23 Okay. And you notice that I'm
24 drawing the electrons in certain places and
25 I'm putting them in pairs. We're going to
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10
Chem. 241 - Lectures 2 and
1 actually find out why I'm doing that, you
2 know, in the next class. But for now, just
3 note as to how the electrons go. They don't
4 go randomly around the atoms; they go in
5 pairs and they're separated by certain
6 angles.
7 * * *

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