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Chemistry: Examining Atomic Structure

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  • Type: Video Tutorial
  • Length: 13:45
  • Media: Video/mp4
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  • Access Period: Unrestricted
  • Download: MP4 (iPod compatible)
  • Size: 148 MB
  • Posted: 01/28/2009

This lesson is part of the following series:

Chemistry: Full Course (303 lessons, $198.00)
Chemistry Review (25 lessons, $49.50)
Chemistry: Atoms, Molecules, and Ions (10 lessons, $16.83)
Chemistry: Atomic Structure (3 lessons, $4.95)

Professor Harman explains isotopes and atomic mass in this lesson covering atomic structure. Most elements exist in nature as more than one isotope. Isotopes are atoms that have the same number of protons as the element but a different number of neutrons. The number of protons always remains the same, as this number (also known as the atomic number) is what determines the element. Prof. Harman also introduces atomic mass units, or amu's, which are a more convenient unit for describing the very small masses of atoms. Next, Professor Harman explains more about the masses of elements. The amu is derived from carbon 12 and is equal to 1.6605 x 10^-27. The relative atomic mass listed on the periodic table of the elements is a weighted average of the masses of the isotopes of an element. You might also observe that the mass of an isotope is less than the sum of masses of its nucleons and electrons. Professor Harman explains the relationship between mass and energy that Einstein discovered, and binding energy.

Taught by Professor Harman, this lesson was selected from a broader, comprehensive course, Chemistry. This course and others are available from Thinkwell, Inc. The full course can be found at http://www.thinkwell.com/student/product/chemistry. The full course covers atoms, molecules and ions, stoichiometry, reactions in aqueous solutions, gases, thermochemistry, Modern Atomic Theory, electron configurations, periodicity, chemical bonding, molecular geometry, bonding theory, oxidation-reduction reactions, condensed phases, solution properties, kinetics, acids and bases, organic reactions, thermodynamics, nuclear chemistry, metals, nonmetals, biochemistry, organic chemistry, and more."

Dean Harman is a professor of chemistry at the University of Virginia, where he has been honored with several teaching awards. He heads Harman Research Group, which specializes in the novel organic transformations made possible by electron-rich metal centers such as Os(II), RE(I), AND W(0). He holds a Ph.D. from Stanford University.

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Atoms, Molecules, and Ions
Atomic Structure
Examining Atomic Structure Page [1 of 3]
With the help of a mass spectrometer, we identified two major isotopes for carbon; carbon-12, meaning 12 total
nucleons, protons and neutrons, and carbon-13, containing one more neutron than the carbon-12. So let’s look a little
bit more carefully as to what we’re talking about here, and we’ll define our terms a little bit.
Remember, first of all, the nucleus has a volume much, much, much smaller than the total volume of the atom, and
inside the nucleus is where we find the protons and the neutrons. The rest of the space occupied by the atom is just
simply electrons. Now, the difference between carbon-12 and carbon-13 we said was one neutron. Notice that the
number of protons in carbon-12 and carbon-13 are exactly the same and, in fact, the number of protons defines those
atoms as being carbon. We oftentimes will describe that in this notation here. You’ll notice the C for carbon. The
twelve above is the mass number. That simply is telling you the total number of nucleons, total number of protons and
neutrons. The lower number, if you see a lower number, sometimes this is not there, is the atomic number, and that
tells you the number of protons. Well, notice that the atomic number for both carbon-12 and carbon-13 is six. It’s
exactly the same. And again, what defines these two guys as carbon is that they have six protons. If we put in
another proton, we get a different atomic number and therefore a different element.
So, in addition to carbon-12 and carbon-13, we also actually have another isotope that is not abundant enough on the
earth’s surface to be significant to us, but it is, in fact, important, because it’s a radioactive element. The nucleus is, in
fact, not stable, and we’ll have lots more to say about that later, but you may have heard of radioactive dating, or
carbon dating, and that has to do with the decomposition of this nucleus. And so we’ll talk more about that a little bit
later.
But right now the point is that all of these isotopes, they all have six protons, identifying them as carbon, and varying
numbers of neutrons. Therefore their mass number changes. When I say then carbon-14, again I’m identifying it as
carbon, six protons, fourteen nucleons total, so therefore it must have eight neutrons in this case.
Now, when we deal with neutrons or protons, they’re really small and it doesn’t make a whole lot of sense, if we’re
talking about the mass of an individual neutron or proton or carbon atom, to deal with a massive unit, like a kilogram.
We know a kilogram is a big, huge macroscopic thing. And so let’s define now a mass unit that’s more appropriate for
the delicate atom or proton. What we’ll do is take a carbon atom, which has a mass of 1.9924 × 10-26 kilograms, and
we’re going to divide it by twelve, because there are twelve nucleons in that carbon-12 atom. Why carbon-12? Why
not something else? This is just a super stable isotope of carbon, and so it’s very convenient to use that as a
standard. But we could have chosen anything we want. But, by convention, scientists have all agreed that that’s
going to be our mass standard. So we’re going to take one carbon-12 atom, divide it by twelve nucleons, six protons,
six neutrons, and that gives us then 1.66 × 10-27 kilograms. Okay, that is our conversion factor. 1 amu we will define
as 1.66 × 10-27 kilogram. So, if we have a mass in kilograms, we know a quick way then to convert it atomic mass
units, or vice versa.
If that’s true, if I have one carbon-12 atom in my hand, it’s going to weigh exactly 12 amu, because we defined it that
way. If I have one hydrogen-1 atom, in other words, hydrogen with just a proton in it, no neutrons, that’s going to have
a mass very, very close to 1 atomic mass unit. Not exactly, as we’ll see why in a little bit. If I have a deuterium atom,
that’s hydrogen with one neutron, so total number of two, a mass number of two, that’s going to have a mass of very,
very close to two atomic mass units. If I have a tritium atom, where the mass number of three, that’s going to be really
close to 3 amu. You get the idea. It makes it much, much easier when we’re talking about atomic mass units,
because that effectively tells us very close to the number of nucleons in the nucleus. That’s good.
When we look at a periodic table then, remember that the different elements are defined by the number of protons
they have. So what makes sulfur sulfur is that it has sixteen protons. It’s got an atomic number of sixteen. What
makes beryllium beryllium is it’s got four protons, an atomic number of four. What makes nitrogen-15 different than
nitrogen-14 is the number of neutrons, just to remind you. It’s still nitrogen, because we haven’t changed the number
of protons in that case. Well, one of the bits of information that you’ll see on a periodic table again is this number at
the top, the atomic number. Once again, just tells us the number of protons. The elemental symbol, which again is
dictated by the number of protons. And then this guy down here, the relative atomic mass. If carbon existed only as
carbon-12, this would be the exact mass of one carbon atom. But, in fact, notice that this number is not twelve. It’s
not twelve because, as you would guys, because I just led you this way, carbon exists as two major isotopes, carbonAtoms,
Molecules, and Ions
Atomic Structure
Examining Atomic Structure Page [2 of 3]
12 and carbon-13, and so this is, in fact, an average mass for a carbon atom, not the actual mass of carbon-12 or
carbon-13. It’s a weighted average of those two, given the fact that there’s a lot more carbon-12 than carbon-13 here
at the earth’s surface.
So let’s define that a little bit better. Again, what I want is a weighted average of those different atom types. So
carbon-12 weighs exactly 12 amu. I multiply that by its abundance point, 98.8. In other words, there’s 98.89%
carbon-12. That’s going to give me this number, and I then multiply carbon-13 by its natural abundance. It gives me
this number. And I just simply sum those two together and that give me a weighted average then of the two. Notice a
weighted average is different than a simple average. If I just took the average of twelve and thirteen, it would be 12.5,
but that would not pay attention to the fact that there’s a heck of a lot more carbon-12 around. So I need to multiply it
by abundances, so that I get a weighted average, the appropriate ratio of these guys. So again, you’ll notice this
number is a little tiny bit bigger than twelve, because again there’s about 1% of carbon-13 at the earth’s surface. And
I’m not worried about carbon-14 here, because it’s just too small of an abundance.
Okay, so if you actually had to do a problem, what it might look like is this, atomic weight of an element corresponds to
the weighted average of all its naturally occurring isotopes. Let’s talk about neon. Neon exists as neon-20, -21, and -
22. The dominant form of neon is neon-20. And again, what separates these guys is the number of neutrons. Okay,
so you want to figure out the atomic mass for neon and you’ve got to take the weighted average of those isotopes.
So, once again, you would be presented typically with the data for the masses of those different isotopes. If you
weren’t presented with them, you could go in the lab and measure them with a mass spectrometer, and then you get
percentages. And again, that’s something that you could get with a mass spectrometer also. You just simply collect
this data and the relative heights on the mass spectrum will tell you the abundances. And again, the mass on the xaxis,
if you remember, that’ll tell you the masses. Multiply abundance by the mass and you end up adding those guys
together, just like we did with carbon, and you end up with a weighted average for neon of 21.18. Okay?
Now, here’s something peculiar. Neon-20 has got a mass of 19.992. I’m going to add the mass of exactly one
neutron to neon-20 and I end up with 21.000 amu. Notice that that number is not the same as the actual experimental
value, 20.993. Now, I know, you could look at that and say, “That’s not a big difference. Maybe somebody just
rounded off wrong or used a Casio calculator, or something like that”. I take it all back – I don’t really mean that. In
this case, the problem is not in the calculation. There is, in fact, a difference between the weight of the neon-20 and
one neutron and the neon-21. And that difference has to do with something you’ll see in just a moment.
I’m going to show you one other example of this, and then we’ll get to what’s going on here. But again, let me just
leave you with the notion that if I take neon-20 and I add a neutron, I don’t get the mass of neon-21, even though we
just said neon-21 is just one more neutron than neon-20. Something’s going on here.
Okay, it becomes even more apparent if we talk about a whole atom. In other words, let’s think of a simple atom,
helium-4. Okay, helium-4 has got two protons, two neutrons, two electrons in it. Well, I could take the mass of a
proton, I know what that is, multiply by two, take the mass of a neutron, multiply by two, the mass of an electron,
multiply by two, add that all together, and the mass that I get is significantly larger than the mass of an actual
authentic helium atom. So something’s going on. There’s 5 × 10-29 kilograms missing. Now, I know, that seems like
a pretty small number, but this is per atom that that is missing. If I had a large collection of helium, that’s a significant
amount of mass that’s just gone.
Now, we love to think about mass being conserved. For the most part, that work’s pretty well. But, in fact, when we’re
talking about changing the nucleus of an atom, mass is, in fact, not conserved. The reason is because of this very
famous relationship that you’ve certainly seen before by Einstein, that energy is equal to mass times speed of light
squared. Now, what that says is that energy and mass are, in fact, interchangeable. And when these guys come
together, there is so much energy released that there is a loss of mass, released again in the form of energy. Well,
put that in reverse. If you took helium-4, if you could take a helium-4 atom and break it apart into protons and
neutrons, you have to put in an enormous amount of energy, in order to do that. And, as a result, the mass actually
would increase and you’d get your particles back here. So, because there is such an enormous difference in energy,
that’s called the binding energy, the energy that holds that nucleus together, and it’s an incredible amount of energy.
So much so that normally we don’t have to worry about the nucleus changing when we’re doing chemical reactions.
We can leave that to the physicists. That’s nuclear chemistry. We’ll actually talk a little bit about that, but not for a
Atoms, Molecules, and Ions
Atomic Structure
Examining Atomic Structure Page [3 of 3]
while. We’re going to just assume that the nucleus is not changing. We’ll go ahead and buy into one of Dalton’s
postulates, that the nucleus is indivisible, even though we know that that’s not completely right. But that’s going to be
the assumption we can make. So, if we do that, then energy and mass will be conserved, as far as we’re concerned,
at least we won’t be able to tell the difference and everything will be okay. It’s only if the nucleus is changing that the
energy differences are so great that the mass will be changing enough that we can actually measure that difference.
So that’ll be important for us to remember when we look at nuclear chemistry, but not when we’re talking about
chemistry in general.
So, okay, where are we? Well, we’ve identified mass number as the total number of nucleons, meaning protons and
neutrons. We’ve identified atomic number as being the number of protons, and the atomic number defining the name
of the element. We’ve also said okay, we’re going to go ahead and assume now that all atoms maintain their nucleus.
We’re not going to worry about nuclear reactions. We’re just going to just consider chemical reactions, in other words,
the rearrangement of atoms to go from some molecules to different molecules.

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