Download Electron Configuration (You will have to read this more than once to

Electron Configuration (You will have to read this more than
once to understand it. There’s no way you can get it just by doing it once).
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OK, so up to now we really haven’t spoken about electrons all that much. We looked at the different
models of the atom up to Rutherford. Democritus thought atoms were just piles of shapes. Dalton said
that all atoms were spheres and that all molecules were combination of spheres. Thomson realized that
those spheres actually had positive and negative pieces. Rutherford’s gold foil experiment showed that
the positive mass isn’t actually spread out over the whole sphere, but actually concentrated into a really
small nucleus. This was a great step forward, but if you remember it left questions.
Question! You know opposites attract. You know electrons have a charge of -1 and that protons have a
charge of +1. You also know that protons only exist in an extremely dense core of an atom called a
nucleus. Finally you know that electrons are outside of that nucleus. So here’s the question: In a
classical view of the world what should happen to the electrons and the protons in the model below?
If you said they should collapse together then great. If you didn’t come up with that you really need to
think a little harder about it. Ok, what happened next was and is a major leap forward. This is where
Bohr’s model comes in. Here is Bohr’s model. You need to remember this for pretty much every test,
quiz, and homework because it’s going to come up over and over and over and over and…
So, what’s the big deal about that? He said electrons couldn’t just be anywhere (actually a guy by the
name of Louis DeBrogalie said that, but Bohr came up with the model of what it would look like)! He
said they have to be in specific places described by their Energy Levels. The first level closest to the
nucleus is level 1. This is often referred to as the first energy shell. The next level is called the second
energy level or second shell. All the shells just go up by 1 number as they go out further and further from
the nucleus. There is no such thing as energy shell 1.1 or 1.2. There are only whole number shells.
If you have ever heard of the term quantum this is what they were talking about. When things are
quantized they can only be of very specific values. Quantized things can not exist in intermediate
values 1.1, 1.2, 2.4, and other non-whole number values are not allowed. This is the first quantized
number that we have encountered in class. It was the first one scientists found too. Because it was the
first quantum number to be discovered it is called the principal quantum number, and it is given the
nickname “n” so that lazy chemists don’t always have to write out principal quantum number. Remember
“n” is talking about what energy level an electron is in. So, in a classical view of the world where can an
electron be compared to the nucleus? In a quantized view where can the electron be compared to the
nucleus (before asking for help look at the two models above and try to figure out the difference).
Ok, so, if electrons get excited they move to higher energy orbits. At some point they lose that energy
and go back to the lower levels. When this happens light is emitted. Given what you know about
quantized energy levels are all colors of light going to be emitted? Or just a select few? Why?
This is where our buddies Schrodinger and Heisenberg come in. They took the idea that the energies that
electrons could hold are quantized and they ran with it. Heisenberg went off and studied the nature of
light and the nature of electrons. As it turns out light is made up of really tiny particles about the same
size and having about the same energy as electrons. Einstein named these small particles of light
photons. Heisenberg realized that if you use light to detect where an electron is you’re going to have a
problem. Since light is the same size and can have about the same energy when you use light to detect an
electron you’re also moving it or changing it’s direction. To get a sense of this we’re gonna imagine that
a photon and an electron are like a cue ball and an 8 ball on a pool table. If you were blindfolded and had
to detect where the 8 ball was you could just roll the cue ball up and down the table a couple of times
until you heard the “clack” of two pool balls colliding. Well, you could hear where it was, and could
have a good guess of what it’s position was, but as soon as they clacked you have no idea where it went
right. Since both of the pool balls are the same size if you use one to detect where the other is you’ll
knock them both off course. Draw a picture of this and add some labels to show me you have an idea
of what’s going on.
Let’s take that pool analogy one step further. Schrodinger agreed with Heisenberg that you couldn’t
positively know where 8-ball was using the cue ball. But he said “ahh who cares, it’s got a 90% chance
of being on the pool table anyway, and only a 10% chance of having ever been knocked off the pool table
and rolling somewhere on the floor.” In a sense, he tried to get away from the limitations of not knowing
the exact location or the direction an electron was heading by saying “Well, it’s pretty darned likely we
can find it here.” To find out what “here” looked like he developed his own set of equations, and what he
figured out is that electrons live in differently shaped orbitals that have different energy levels. The very
first orbital looked just like a sphere. He called it the 1s orbital. Please draw a sphere and label it 1s.
To get to this first orbital Schrodinger didn’t just use the principal quantum number. In fact he had to use
4 quantized numbers to describe the shape of the orbital that housed the electron. The second quantized
number is called the “angular momentum quantum number.” This number describes the shape of an
electron orbital. In that sphere above the 1 is talking about the principal quantum number, which is the
energy shell. The “s” is describing the shape. You can think of “s” as meaning “sphere.” Now there are
rules that go along with the angular momentum quantum number. The first rule is that chemists gave it
the nickname “l” (as in Lazy). Remember that the principal quantum number, or n, as we chemists call it
can be any whole number larger than zero. Well, “l” can be 0 as well as all positive numbers equal to or
less than n-1. If n=1 what numbers can “l” be?
If n=2 what numbers can ”l” be?
If n=3 what numbers can “l” be?
As it turns out chemists assign specific letters to each one of those values of l. If l=0 they call it s (as in
sphere). If l=1 they call it “p.” I’ll refer to it as a “p-orbital.” If l=2 they call it “d.” I’ll refer to it as a
“d-orbital.” If l=3 they call it “f.” I’ll refer to that as an “f-orbital.” Now theoretically a lot more orbitals
exist, but chemists don’t care because you’ll almost always find an electron in a s, p, d or f orbital. sorbitals look like spheres, p-orbitals look like bowties, most of the d-orbitals look like 4 leaf clovers.
Don’t worry about what the f-orbitals look like. They’re crazy.
Chemists started out happy with just these two numbers, but then they started testing the orbitals that they
came up with in magnetic fields and they realized something was wrong. As it turned out some of the
orbitals could have very specific orientation. For instance, the p-orbital (the one that looks like a bow tie)
can actually go up and down, left to right, or front to back. Draw 3 lines and label them X, Y, and Z
just like a 3-D Cartesian Coordinate system. Do that 3 times. And on each one put a p-orbital that
is going either up and down, left to right, or front and back. Label them 2px, 2py and 2pz.
Depending on which axis they are on.
As it turns out the fact that orbitals lined up in a magnetic field in different ways could also be quantized.
This number is called the “magnetic quantum number.” Lazy chemists just call it “m” for magnetic.
As it turns out “m” can any whole number value between –l and l including zero. So, let’s say l=3, what
values of m are allowed?
The final quantized value is the “Spin quantum number” which only has 2 values +1/2 and -1/2. This
basically means that an electron can spin clockwise or counterclockwise.
Let’s summarize this stuff into one big table that you’ll probably need to refer back to a bunch of times.
The
values
“n” can
equal
1
2
2
3
3
3
4
4
4
4
The
values
“l” can
equal.
0
0
1
0
1
2
0
1
2
3
Name
of
orbital
given
to “l
value”
s
s
p
s
p
d
s
p
d
f
“m,” which is
equal to whole
number values of
–l to l. inc. 0.
0
0
-1, 0, 1
0
-1, 0, 1
-2, -1, 0, 1, 2
0
-1, 0, 1
-2, -1, 0, 1, 2
-3, -2, -1, 0, 1, 2, 3
Number
of
acceptable
values for
“m.” Or
number of
orbitals
available
to a given
“n”
Total
number of
orbitals
available for
a given
value of “n.”
Each orbital can
hold 2 electrons:
one that is spinning
clockwise and one
that is spinning
counter-clockwise.
This is how many
electrons the “n”
shell can hold.
1
1
3
1
3
5
1
3
5
7
1 for n=1
4 for n=2
2e
8e
9 for n=3
18e
16 for n=4
32e
Energy
order.
Lowest
energy
orbitals
get a
value of
1.
1
2
3
4
5
7
6
8
10
13
There are 3 rules for putting electrons into orbitals. “Aufbau principle,” “Pauli exclusion principle,”
and “Hund’s rule.” Aufbau means “building up.” It means you put electrons in the lowest energy
orbitals first. So, n=1 orbitals first, then n=2 and so on. The Paulie exclusion principle means no 2
electrons can have the same 4 quantum numbers. So, you can’t put two electrons in the exact same place
the exact same way. Hund’s rule means you put 1 electron spinning clockwise (just draw an arrow
pointing up) in each empty orbital of the same energy before you start putting electrons in spinning
counterclockwise (draw the arrow facing down). So, you’d put the first electron pointing up (aka spin up)
in the 1s orbital and then you’d put one pointing down (spin down) in the 1s orbital. Then put them in the
2s and then in the 2p, but start putting all of them in spin up, then pair them up with a spin down electron.
Mrs. Webber has a great worksheet called “Atom City Electron Hotel.” Let’s use that and get practice.