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Paramagnetism and diamagnetism

Created by Jay.

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  • blobby green style avatar for user Jason Allen
    Is there a difference in the paramagnetism value/effect between those elements like Cl that are exhibiting paramagnetism only because of the final unfilled sub-shell (3p in this case) in the p-orbital?

    In comparison to say Cr or Cu which have more sub-shells only partially filled and hence all 4s and 3d spins in the same direction?
    (32 votes)
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  • blobby green style avatar for user Justin Rider
    I have a question, why is Mg and Ca paramagnetic even though they have paired electrons in their s orbitals?
    (11 votes)
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    • leafers sapling style avatar for user Marc Johnson
      That is a good question, but its answer comes in two parts, and the second part requires some information you haven't encountered yet in order to answer fully. Ready?

      OK, here we go: (Part 1:) A single (isolated) Mg or Ca atom in its electronic ground state is diamagnetic, as you would predict from its ground-state electron configuration of (1s)2(2s)2(2p)6(3s)2.

      (Part 2:) However, a chunk of Mg or Ca metal contains a lot of Mg (or Ca) atoms. You haven't yet learnt about how the atomic orbitals (AOs) on one atom interact with the AOs on another atom to make molecular orbitals (MOs), but they do when the atoms are close enough -- and that's what leads to bonding (or not, depending on the interactions). It turns out that some of the interactions between the 3p orbitals on different Mg atoms can lead to molecular orbitals that are lower-energy overall than some of the 3s-only combinations, with the effect that some of the electrons in a chunk of Mg metal end up in 3p-based MOs, and can then have the same spin as some electrons in the 3s-based orbitals, for overall paramagnetism.
      We say that the 3s band of MOs overlaps the 3p band of MOs, leading to a combined band that is only partially filled by the valence electrons.
      If this doesn't make much sense yet, don't worry -- it will. Just wait until you get to Linear Combinations of Atomic Orbitals - Molecular Orbital (LCAO-MO) theory, then come back to this question.
      (8 votes)
  • blobby green style avatar for user Otte de Boer
    So, does that mean when e.g. the 3p orbit Al, Si, and P are up-spin, and S, Cl, and Argon are downspin? It does not conform to your definition, but otherwise I cannot see how up-spin and downspin would work. I understand the rest, just wondering when to decide wether it i s +1/2/-1/2.

    Thank you,
    (4 votes)
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    • aqualine tree style avatar for user anderocketech
      All unpaired electrons are labeled spin up from what I recall. It's just convention - it has nothing to do with how orbitals really work. By convention, you fill up all of the sub-shells with 1 electron first (meaning they are all spin-up) before adding a second electron to each sub-shell. If we represent the spin as + and -, we can say that the Chlorine outer p-orbital is filled like this: (+,-), (+,-), (+, ) whereas the Silicon outer p-orbital is filled like this: (+, ), (+, ), ( , )
      (16 votes)
  • purple pi purple style avatar for user brewbooks
    I have read that hemoglobin is paramagnetic when it is deoxygenated and diamagnetic when it has oygen bound. Can someone help me understand what is going on in a simple manner. One explanation I read is that "Hemoglobin without bound oxygen molecules, deoxyhemoglobin, is paramagnetic because of the high spin state (S = 2) of the heme iron. In contrast, oxygen-bound hemoglobin, oxyhemoglobin, has low spin (S = 0) and is diamagnetic (Pauling &Coryl 1936). " I know the iron in heme is Fe2+
    PS - The source if my info on Hemoglobin was http://www.scholarpedia.org/article/Functional_magnetic_resonance_imaging
    (4 votes)
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    • spunky sam blue style avatar for user Ernest Zinck
      The Fe²⁺ ions in deoxyhemoglobin are coordinated to five N atoms. In oxyhemoglobin, the Fe³⁺ ions also have an O₂ molecule.
      According to crystal field theory, these surrounding ligands split the orbitals into two groups, with three of them being slightly lower in energy than the other two.

      Fe²⁺ is a d⁶ ion. The electrons could be arranged either as d²d¹d¹ d¹d¹ (high spin) or d²d²d² d⁰d⁰ (low spin). The high-spin state has S = 4×½ = 2. The Fe²⁺ in deoxyhemoglobin is high-spin.
      Fe³⁺ is a d⁵ ion. The electrons could be arranged either as d¹d¹d¹ d¹d¹ (high spin) or d²d²d¹ d⁰d⁰ (low spin). The low-spin state has S = 1×½ = ½. The Fe³⁺ in oxyhemoglobin is low-spin.

      However, oxyhemoglobin is diamagnetic (S = 0).
      O₂ is a diradical.
      The reaction when oxygen binds to hemoglobin is
      Fe²⁺ + O₂ → Fe³⁺ + O₂·⁻ (Fe is oxidized; O₂ is reduced)
      The O₂·⁻ still has one unpaired electron.
      When low-spin Fe³⁺ binds to O₂·⁻, the spins of the two unpaired electrons couple strongly
      (S = ½ - ½ = 0), so oxyhemoglobin is diamagnetic.
      (7 votes)
  • duskpin ultimate style avatar for user MS17155 - Shivanshu Siyanwal
    Why does a moving charge produces magnetic field around itself?
    (4 votes)
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  • spunky sam blue style avatar for user Ayan Gangopadhyay
    I don't get how the diamagnetic substances are repelled by the magnetic field. Can anyone help me out? Any help is much appreciated.
    (5 votes)
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  • piceratops tree style avatar for user phoskere
    There is a another category i have studied, it is called ferromagnetic, what is it?
    (3 votes)
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    • old spice man green style avatar for user Matt B
      Nice observation! In fact, there is yet another type: antiferromagnetism (and some even consider ferrimagnetism to be another, separate category).
      Ferromagnetic simply means that a substance is attracted very strongly to a metal. The common ferromagnetic metals are iron, nickel, cobalt and most of their alloys, some compounds of rare earth metals, and a few naturally-occurring minerals such as lodestone.
      (5 votes)
  • blobby green style avatar for user eleonoramatic139
    I think the question was already asked here before, but I don't think anyone answered the part I'm thinking of. So how do you determine if a compound is diamagnetic or paramagnetic? I assume it's the same principle as with atoms and ions, but I don't quite understand how can you see that from the Lewis structure or the number of valence electrons. For example the O2 molecule is paramagnetic and even has an even number of valence electrons and it also doesn't have unpaired electrons that you can see on a Lewis structure. Any help is appreciated, thanks!
    (2 votes)
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  • blobby green style avatar for user punhal.rajper
    I have been through some examples, which ended with the O2 molecule being paramagnetic. While N2 is diamagnetic in nature, even though both satisfy their octet rule as well as orbital bonding(pairing of electrons).
    Please! help me out
    (3 votes)
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    • leaf red style avatar for user Richard
      So in this situation it's important to be aware of the different bonding theories in chemistry. The two main ones are valence bond theory and molecular orbital theory. Essentially valence bond theory, which is used in this video, views bonding electrons as existing between only the two bonding atoms. Molecular orbital theory on the other hand views bonding electrons as existing over the entire molecule. Valence bond theory makes frequent use of hybrid orbitals while molecular orbital theory uses bonding, nonbonding, and antibonding orbitals. Both theories have their usefulness in explaining certain phenomena, but also have their limitations. Explaining why a molecule of N2 is diamagnetic, but O2 is paramagnetic is best done through molecular orbital theory.

      Using molecular orbital theory, or MO theory, we can construct energy diagrams similar to those in valence bond theory, but are populated by various molecular orbitals. We fill them similarly using the valence electrons with the aufbau principal and Hund's rule, but just into MO's which include bonding, nonbonding and antibonding orbitals.

      Using the MO diagram for N2, we fill into it the 10 valence electrons from the two nitrogens. We find that all the electrons are spin paired and that it is diamagnetic. We can also predict that the bond will be a triple bond formed from a sigma bond and two pi bonds from the MO diagram.

      The MO diagram for O2 is slightly different from N2, but we fill the 12 valence electrons from the two oxygens in the same manner. When we do this we'll have two unpaired electrons in the antibonding pi orbitals. These electrons account for O2's paramagnetism. The MO diagram also predicts the bond order to be 2 meaning it'll have a double bond which is what we observe.

      So chemical bonding can be complex and certain bonding theories may be insufficient to explain certain phenomena. So there isn't a single correct bonding theory and we have to be aware of different theories to have a more accurate understanding of chemistry.

      Hope that helps.
      (2 votes)
  • piceratops ultimate style avatar for user Braden Cummings
    So is N paramagnetic with its 3 electrons in 2p with up spin. I know that 2 electrons with parallel spin makes is paramagnetic but what about 3?
    (3 votes)
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Video transcript

- [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.