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Current time:0:00Total duration:8:05

- [Voiceover] We've already seen that the allowed values for
the spin quantum number are positive one half
and negative one half so an electron can have spin up or an electron could have spin down. And remember spin is in quotation marks because we can't really visualize an electron spinning on its axis. That's not really what it's doing. So we just called it
the spin quantum number. And so let's say we have. Let's say we have two electrons and each of our electrons has spin up. So lemme see if I can
draw that situation here. So we have two electrons with spin up. Well an electron is a moving charge. Moving charges produce magnetic fields. So an electron is really
just a tiny magnet. And when you have two
electrons with parallel spins, the magnetic fields of those electrons add together. So we call the situation paramagnetic. So this situation here is paramagnetic. The magnetic fields of the
electrons add together. If you have a situation
where you have one electron with spin up and one
electron with spin down, the magnetic fields of those electrons cancel each other out. And so we call this situation diamagnetic. And so let's get some better definitions for paramagnetic and diamagnetic. So let's move down to here. And let's look at the
definition for paramagnetic. So something that's paramagnetic has one or more unpaired electrons. So we talked about an example where we had two unpaired electrons. But of course you could just
have one unpaired electron. Right so that's like a tiny magnet with its own magnetic field. And so something that's paramagnetic is pulled into an external magnetic field. It's attracted to an
external magnetic field. And we can figure out if a sample is paramagnetic or not by
using this special balance that I have. I have this picture of this
balance drawn down here. So let's say that our paramagnetic sample is in here. So right there in magenta. And we haven't turned on the magnet yet. So here we have a magnet. There's a north pole and a south pole. So before we turn the magnet on, let's just say that
our paramagnetic sample is balanced by some
balancing weight over here on the right side. Right so there's a pivot point right here but we have everything balanced perfectly. Alright so let's now turn the magnet on. So we turn the magnet on
and the magnetic field lines go from north pole to
south pole like that. And if we have a paramagnetic sample. With one or more unpaired electrons, our paramagnetic sample is pulled into this external magnetic field
that we've just turned on. And so this is pulled down, right? So this whole part is pulled down. And so let me go ahead and redraw it here. And so this would be pulled down into the magnetic field and so our paramagnetic sample is
pulled into the magnetic field. Right what does that do to our balance? Well of course that's going
to pull this side down. And so that's going to pull and our balance is going to
rotate about this axis, right? And so this part's gonna go up. So just simple physics. So this weight's gonna go up. It's like our paramagnetic
sample has gained weight. And of course it hasn't gained weight, just experiencing a force. There's a magnetic force because it is a paramagnetic substance. And so this balance
allows us to figure out if something is paramagnetic or not. Let's look at the
definition for diamagnetic. So for diamagnetic all
electrons are paired. So we have, if we have
spin up, we have spin down. And so the magnetic fields cancel. And so a diamagnetic sample
would not be attracted to an external magnetic field. Actually it produces
its own magnetic field in the opposite direction. So it's actually weakly repelled by an external magnetic field. So we have these two definitions. Paramagnetic and diamagnetic. And we can figure out if atoms or ions are paramagnetic or diamagnetic by writing electron configurations. So let's look at a shortened version of the periodic table. And let's look at some elements. And let's figure out
whether those elements are para- or diamagnetic. Let's start with helium. So helium right here. We need to write the electron
configuration for helium. So this would be 1s1 and then we get 1s2. So I'm assuming you already know how to write your electron configurations. So we have 1s2 which means we have two electrons in a 1s orbital. Here's our 1s orbital. We have two electrons and
they must be spin paired. Right so the electrons
are completely paired and that means that helium is diamagnetic. Helium is diamagnetic. So helium atoms I should say. Let's do carbon next. Let's find carbon. Let me change colors here. Here's carbon on the periodic table. If I wanted to write an electron
configuration for carbon, well it would be 1s2. Right so I'll start 1s2. Then we have 2s2. So 2s2. And then we have, we're in the 2p1 and then 2p2. So 1s2, 2s2, 2p2 is the
electron configuration for carbon. If you write in orbital notation. Right so we would have
our 1s orbital here. And our 2s orbital here. And then we have three
2p orbitals like that. So we'll put in your electrons. We have six electrons. Alright so two in the 1s orbital. So we put those in. Two in the 2s orbital. We put those in. And remember Hund's rule, right? We have two electrons in the p orbital. But we don't pair those spins, right? We don't pair those spins. And so we have. We have unpaired electrons. We have unpaired electrons here for carbon when we draw out the orbital notation. And unpaired electrons means that carbon is paramagnetic. So carbon is paramagnetic. Carbon atoms anyway. Let's do sodium next. So let's find sodium down here. So here's sodium. We need to write the
electron configuration. Right, so that would be 1s2. So let's write 1s2 here. 2s2, and then we have 2p6. So 2p1, 2p2, 2p3, 2p4, 2p5, 2p6. So 2p6. That takes us to the 3s orbital. Right so one electron in the 3s orbital. So 3s1. So 1s2, 2s2, 2p6, 3s1 is the electron configuration for sodium. If we did that on our
orbital notation, right? We would have 1s orbital. Alright so we have two
electrons in the 1s orbital. 2s orbital, we have two
electrons in the 2s orbital. 2p orbitals, right. We have one, two, three, four, five, six. And then we have 3s1. Right so we have the
3s orbital right here. One electron in the 3s orbital. We'll notice one unpaired electron. An unpaired electron means paramagnetic. So sodium. Sodium is paramagnetic. Sodium atom anyway. Finally let's do sodium ion. So Na+. So the sodium atom has equal numbers of protons and electrons. But the sodium ion, we've
lost one of those electrons. Right so we're going to lose
this outer electron here. Right so the sodium ion has this for an electron configuration. 1s2, 2s2, 2p6. And so we lose this one electron. Notice for the ion now we
have all paired electrons. Right so everything here is paired. And if you have all paired electrons, we're talking about diamagnetic. So while the sodium atom is paramagnetic, the sodium, I misspelled that. The sodium ion is diamagnetic. And so it's just about writing your electron configurations and thinking about the definitions for paramagnetic and diamagnetic.