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My issue

I have a significant problem with the following excerpt from a general chemistry text:

Note that ON stands for oxidation number (state). The excerpt is with regard to polyatomic anions - specifically, chloride ion and its related polyatomic anions.

It now follows that as oxygen-atom content increases by one, ON for the atom of variable ON increases by two because each added O has an ON of -2 and O is more EN than the atom of variable ON - thus, bonding O to this atom makes the atom more electropositive.


If electropositivity is defined as the ability to stabilize positive charge, how does a highly partially positive central chlorine atom have this ability? How can EP increase?


If electropositivity is more formally defined as the ability for an atom to give up electrons within a covalent framework (note that this is the converse of the definition of electronegativity), how exactly does a highly oxidized chlorine atom, such as the chlorine atom in the perchloric ion, have the ability to give up any more electrons? It already has no valence electrons to give up if we stick with an ionic model to assign oxidation states!


If we look at this intuitively, and agree that more electropositive elements are more likely to bear positive charges - i.e. sodium (Na) is highly electropositive and it is very easy for Na to be stripped of its one electron to become isoelectronic with a noble gas - then why would chlorine become more electropositive as it is further oxidized?

If we consider the successive ionization energies of chlorine, they become more and more massive! So how in the world would chlorine be more electropositive (willing to give up electrons within a covalent framework) as it is further oxidized?

(I do recognize that ionization energy of an electron is not a measure of electronegativity but still it is highly correlated with EN).


Professor's response

My professor, respectfully, didn't offer much a defense. He said something about how positive charge doesn't exist because it's the lack of valence electrons rather than the presence of something material. This, to me, is sidestepping the argument.

He also asked me which would be more likely to bear a positive charge - Cl or O? Obviously Cl; it has electrons way further from the nucleus than oxygen, so of course Cl would more easily be stripped of an electron than O. However, this again, I believe, is a mistake, because we cannot consider electronegativities of elements and expect them to hold within ions and molecules. Carbon, for example, isn't very electronegative at all. But no one doubts the electron-withdrawing powers of a carbonyl carbon, or better yet, the electron-stabilizing powers of a ... CARBOCATION.

Further, last semester, I distinctly remember correcting him with regard to this electronegativity business; he said that a more negative ion would be more electronegative. I had to correct that statement; a more negative ion would be less able to stabilize valence electron density and less able to attract valence electrons if bound within a covalent framework, so the correct term would be instead more electropositive.

Does anyone here agree with me? Sometimes I feel that it's not the argument that matters as much as your tenacity and ability to present the argument. I lacked both tenacity and presentation skills when presenting the above argument.

I came close to telling him that his entire take of electronegativity was fundamentally flawed when I told him that "I don't associate a positive charge with electropositivity and a negative charge with electronegativity" but that was met with nothing to the contrary (which would have been welcome!) - because no, positive charge has nothing to do with electropositivity; in fact, it suggests the opposite of electropositivity!

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  • $\begingroup$ They must mean the central atom becomes more positively charged, not more electropositive. That atom becomes more electronegative! $\endgroup$ – jerepierre Sep 12 '15 at 1:23
  • $\begingroup$ They do mean the central atom but they insist it's more EP for some reason!!! $\endgroup$ – Dissenter Sep 12 '15 at 2:03
  • $\begingroup$ They are clearly not using the term correctly, or at least in the same way I understand it. Don't get up on it. $\endgroup$ – jerepierre Sep 12 '15 at 3:57
  • $\begingroup$ How do you define EP? $\endgroup$ – Dissenter Sep 12 '15 at 4:13
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There are two definitions of electropositive:

1: charged with positive electricity

2: having a tendency to release electrons

The author is using the first definition, while you are thinking of definition 2, which is the one that usually appears in chemistry books.

Instead of "makes the atom more electropositive" it could say "makes the atom more electrically-positive".

For example according to the calculation data from Paul J. Schupf Computational Chemistry Lab

Starting from a -1 charge for Cl-

Adding one oxygen atom (ClO-) increases the charge on the Cl atom to -0.298.

Adding two oxygen atoms (OClO-) increases the charge on the Cl atom to +0.212.

Adding three oxygen atoms (chlorate) increases the charge on the Cl atom to +0.510

Adding four oxygen atoms (perchlorate) increases the charge on the Cl atom to +0.907.

Another example of a different author using the word "electropositivity" in the same way, but who makes the meaning explicitly clear is Interaction of indole derivatives with biologically important aromatic compounds. Part 22. Importance of simultaneous co-operation of hydrogen-bond pairing and stacking interactions for recognition of guanine base by a peptide: X-ray crystal analysis of 7-methylguanosine-5′-phosphate–tryptophanylglutamic acid complex

MNDO calculations showed that the protonation of the guanine N(7) atom increases the electropositivity of the 8-H atom [ -0.014 e for guanine and 0.079 e for N(7)-protonated guanine],

(bracketed text in original document)

Specifically regarding chlorine, the term "electropositive chlorine" has been used to refer to Cl+, for example in the Canadian Journal of Research

Another example to look at is from the thesis Synthesis of functionalised silanes for use in the asymmetric allylation reaction:

In silicon based C-C bond forming reactions the important factor key to their success is the ability of the coordination number of the silicon atom to be varied. As the number of bonds to silicon increases the electropositivity and consequently Lewis acidity on the atom increases.[reference 1]

Upon expansion (Scheme 3; 11 to 13), and ensuing reduction in the s-character orbital composition at the silicon centre, the electron density decreases.

Therefore the electropositivity and Lewis acidity at the silicon centre is increased. These tetravalent and pentavalent Lewis acidic silianes have been exploited in several Lewis acid-catalysed transformations. Once the shell has expanded to incorporate six subsituents it is unlikely for any further extension of the valence shell to occur.[reference 2]

Silicon becomes more positively polarised with the addition of each ligand. As this occurs there is a shift in electron density. While the electron density is increasing at the ligand, it is decreasing at the silicon centre. The magnitude of this polarization is also dependent on the electronegativity of these ligands.

For a very old example, showing the original meaning of "electropositive" to be the first definition, see William Henry's 1823 The Elements of Experimental Chemistry, Volume 1 where on page 229 there is a section "Electro-positive Bodies" which defines electropositive bodies as bodies attracted to a negative electrode and whose inherent electrical state is therefore positive.

Also, there is an historically-interesting article "Electropositivity" and "electronegativity" J. Chem. Educ., 1938, 15 (6), p 291 discussing the inconsistent usage of the term electropositive and advocating for saying "electron donor" instead (if meaning #2 is intended) to avoid confusion.

There is also a 1949 textbook General Chemistry which uses both meanings in the same sentence, using the older meaning to define the newer meaning:

The electropositive character (tendency to go to a more electropositive condition) of elements, in general, decreases across the periods

Overall, it is confusing to use the word "electropositive" contrary to the definition that usually appears in recent chemistry books, but not really wrong.

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    $\begingroup$ Right, looks like they are confusing 'electropositive' and 'Lewis acid' or 'electrophilic'. $\endgroup$ – julien Apr 9 '16 at 12:23
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I would say that is fundamentally incorrect to say that the addition of oxygens to an atom makes it more electropositive. Electropositivity is the measure of an atoms ability to donate electrons, and the more electrons that are withdrawn by oxygen atoms, the harder it will be for the central atom to donate more electrons to form higher oxidation states. I think what it may have been trying to explain is that a partial positive charge (denoted $\delta +$) is induced in the central atom with to the addition of oxygen atoms.

So to answer your question of “how can electropositivity increase,” I would say that this is done in the opposite manner to how your text proposes, that is by removing oxygen atoms. By decreasing the partial positive charge, the central atoms affinity for its own electrons decreases, and the easier it becomes to share those atoms in a covalent bond.

You also seem to be mixing different ideas when you talk about positive charge. There is formal charge, and there is the oxidation state. Formal charge is a number assigned to an atom in a molecule that tells you how many more pairs of electrons the atom is in control of than is usual for that molecule, and oxidation state tells you the theoretical charge an atom would have if all its electrons participating in bonding were considered completely donated (as in an ionic bond).

Interestingly enough, atoms that are good at accepting a positive formal charge are markedly bad at achieving high oxidation states. Take the molecule 3-chloropropan-1-ol, for instance. In an acidic solution, which atom is more likely to donate an electron pair to share with a $\ce{H+}$ ion? Oxygen, because it is more basic (as a result of its high electronegativity). Now consider the perchlorate ion. As you said yourself, chlorine’s valence electrons are more shielded that oxygens, and they are more available to be shared in a covalent bond with an electronegative element like oxygen, even to the extent as to give Cl(+VII). Just try to imagine $\ce{O5}$. It just doesn’t work! Oxygen is far too electronegative.

Also I am not sure in what way your teacher meant “positive charge doesn’t exist.” What then is a proton?

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  • $\begingroup$ I think he meant that in the same way people say cold doesn't exist $\endgroup$ – Dissenter Apr 8 '16 at 4:51

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