Is the covalent character of sodium chloride affected by its polarity?

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Discussion Overview

The discussion revolves around the covalent character of sodium chloride (NaCl) and how it relates to its polarity and dipole moment. Participants explore the nature of ionic bonds, the calculations of ionic character, and the implications of electronegativity differences in ionic compounds.

Discussion Character

  • Debate/contested
  • Mathematical reasoning
  • Conceptual clarification

Main Points Raised

  • One participant calculates the dipole moment of NaCl and questions the resulting value of 2/3 of an electron, expecting it to be closer to a full elementary charge.
  • Another participant suggests that the electron can spend time near the Na+ ion, indicating a degree of covalent character.
  • It is noted that no bond is 100% ionic, and participants discuss the implications of this on the characterization of NaCl.
  • One participant argues that the solid-state separation used in calculations is incorrect and suggests using the gas-phase value instead, which would yield a different result.
  • Concerns are raised about the point charge approximation for the Na and Cl ions, particularly the "fluffy" nature of the Cl ion affecting calculations.
  • Participants debate whether a bond between the most electronegative and least electronegative atoms could be considered 100% ionic, with differing views on the nature of ionic character in compounds.
  • There is a discussion about the variability of covalent character in ionic compounds and whether it can be assumed to disappear at certain electronegativity differences.

Areas of Agreement / Disagreement

Participants express differing views on the degree of ionic versus covalent character in NaCl and other ionic compounds. There is no consensus on the exact nature of this relationship or the implications of electronegativity differences.

Contextual Notes

Participants highlight limitations in their calculations, including the choice of bond length and the assumptions made about ionic character. The discussion remains open regarding the interpretation of ionic versus covalent character in various compounds.

Tiiba
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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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Well, because it's an ionic bond.
 
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.
 
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?
 
Link tells the story: closer than 0.236 nm and the core electrons start pushing back
 
You also have to take into account that the sodium ion polarizes the chloride ion. This will reduce the dipole moment considerably.
 
Tiiba said:
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.

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.
 
  • #10
Borek said:
No bond is 100% ionic.
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?
 
  • #11
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).
 
  • #12
Borek said:
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).
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?
 
  • #13
Comeback City said:
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?
Probably both dissociate into hydrated ions in water. Seriously, where does this difference matter?
 
  • #14
Comeback City said:
By saying "all others we know", do you mean all other ionic compounds in general?

Yes.

If so, are you saying that there would be no difference between a 60% ionic compound, and say, an 80% ionic compound?

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.
 
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  • #15
Borek said:
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.
That way of thinking makes more sense. Thanks.
 

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