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Polarity of sodium chloride

  1. Jan 16, 2017 #1
    I found out that sodium chloride has a dipole moment of 9 debye, and a sodium-chlorine distance of .28 nm. When I divide one by the other, I get 2/3 of an electron.

    Did my math go wrong somewhere, or is this supposed to happen? I expected something close to a full elementary charge.
     
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  3. Jan 16, 2017 #2

    BvU

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    Last edited: Jan 16, 2017
  4. Jan 16, 2017 #3
    Well, because it's an ionic bond.
     
  5. Jan 16, 2017 #4

    BvU

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    Doesn't mean the electron can't spend a fraction of its time in the neighbourhood of the Na+ ion ?

    Link claims the 0.28 nm is in the crystalline state and somewhat greater than in the gaseous state.
     
  6. Jan 17, 2017 #5

    Borek

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  7. Jan 17, 2017 #6
    I know that, but Ithought it would be at least 90% ionic, given that this is THE classic ionic compound. I mosly wanted to know if my math is right.

    So the electron really spends that much time in the sodium?
     
  8. Jan 17, 2017 #7

    BvU

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    Link tells the story: closer than 0.236 nm and the core electrons start pushing back
     
  9. Jan 18, 2017 #8

    DrDu

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    You also have to take into account that the sodium ion polarizes the chloride ion. This will reduce the dipole moment considerably.
     
  10. Jan 19, 2017 #9
    There are two problems in your analysis.

    1- First you used the solid state NaCl separation in your calculation. Instead you should use the gas phase value , 0.236 nm. If you do this you get ~ 0.79 e. In the solid state, the net dipole moment of NaCl crystal is zero due to the centrosymmetry of the structure. This is , of course, not the case for the gas phase molecule.

    2- Even the value of 0.79 e is not a good metric for the ionicity of NaCl molecule because it is calculated assuming that both Na and Cl are point charges. While this point charge approximation is reasonable for the Na ion, it is not for the Cl ion. The latter is sort of a "fluffy" ion which makes it difficult to represent it as a point charge.
     
  11. Jan 20, 2017 #10
    Wouldn't a bond of the most electronegative and least electronegative atoms (so Fluorine and Francium) be 100% ionic, based on the 0-3.3 scale of electronegativity difference?
     
  12. Jan 20, 2017 #11

    Borek

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    Why should it?

    Sure, it is a best candidate we can think of, but there is no reason to think this particular one will be different from all others we know (of which none is 100% ionic).
     
  13. Jan 20, 2017 #12
    By saying "all others we know", do you mean all other ionic compounds in general? If so, are you saying that there would be no difference between a 60% ionic compound, and say, an 80% ionic compound?
     
  14. Jan 20, 2017 #13

    DrDu

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    Probably both dissociate into hydrated ions in water. Seriously, where does this difference matter?
     
  15. Jan 20, 2017 #14

    Borek

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    Yes.

    No. What I am saying is that they all contain some covalent character. The degree to which they are covalent changes, but there are no reasons to assume it will disappear at some particular value of electronegativity difference.
     
  16. Jan 20, 2017 #15
    That way of thinking makes more sense. Thanks.
     
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