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  1. My textbook says that a $\ce{C^{4+}}$ cation cannot be formed because it requires a lot of energy to remove 4 electrons. Formation of ionic bonds involve "removing" electrons and there seems to be enough energy there. So what's different for carbon?

  2. The textbook also mentions that a $\ce{C^{4-}}$ anion cannot be formed because 6 protons cannot hold on to 10 electrons. Elements like chlorine form ionic bonds and end up with 18 electrons and 17 protons. I know there's just one more electron. So is $\ce{NaCl}$ possible because it's not very hard to hold on to 1 extra electron? In that case at what exact point does it become hard to hold onto the extra electrons?

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    $\begingroup$ Do you know about carbides? Some of them formally contain $\ce{C^{4-}}$. $\endgroup$ – Philipp Oct 13 '14 at 17:23
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$\ce{C^{4+}}$ ions:

Single ions are really only observed in the gas phase. There is absolutely nothing that prevents a C4+ ion from being generated in the gas phase, given sufficient energy. Each successive electron removal requires additional energy though, so by the time you get to 4 electrons removed, you're looking at quite a lot of energy required. Note that this does not mean that $\ce{C^{4+}}$ ions are seen in compounds, just that they can be made in the gas phase. Any cation you can imagine can be produced in gas phase.

The difficulty arises when we want to do something with that carbon cation. If I want to make a compound with it, it can bond with other substances and make new bonds. Once that is done we can evaluate those compounds to tell whether the bond is truly ionic or not. In condensed phases or in solution, we look at how the parts of a compound interact with the particles surrounding them to decide whether to treat them as ions or not.

The best chance at making a C4+ ion would be to bond it to fluorine, which is an extremely electron-hungry element. When I actually try this, I find that the C-F bond is very polar, but the $\ce{CF_4}$ molecule that forms behaves as a covalent compound normally would. For example, it does not dissociate into ions when dissolved or melted and it doesn't form an ionic crystal in its solid form. As much as fluorine wants electrons, it doesn't want them enough to completely steal 4 of them from carbon, because the energy required to remove each electron becomes greater than the energy required to remove the one before it. It never gets to the point of $\ce{C^{4+}}$ and $\ce{F^{-}}$ ions being bonded.

Lead can form a +4 ion, $\ce{Pb^{4+}}$, in compounds such as lead fluoride. The reason that this can happen but carbon cations cannot is that the outermost electrons in a lead atom are much farther from the nucleus than those of a carbon atom and are easier to remove as a result.

$\ce{C^{4-}}$ ions:

In gas form this does not happen. Carbon's first electron affinity is positive, so we would expect to see it form $\ce{C-}$ ions in the gas phase if given free electrons to attract. Adding a second electron to this substance would require it to attract an electron against its already negative overall charge, and that is not energetically favorable.

In order to make the $\ce{C^{4-}}$ion in a compound, a species would have to donate 4 electrons to carbon. The attraction of carbon's nucleus to other atoms' electrons is just too low for that to happen. The best candidates for a cation in an ionic compound of this type would be either lithium or cesium, however, when we try to actually make this compound we get another class of ions, the carbides, which feature $\ce{C2^{2-}}$ ions. In short, it doesn't happen.

Phillip's comment regarding the carbides is a good one. There are a few metal carbides that feature carbon atoms bonding to a metal in the ratios we would expect if it were purely ionic, like $\ce{Mg2C}$. In these compounds the carbon does have a formal charge of -4, but the properties of the substance are such that it behaves like a covalent network compound rather than an ionic compound.

What is the limit of this process? Anions with -3 charge, like nitride or phosphide, do exist, though they require a very weakly electronegative species, like an alkali metal or alkaline earth metal, to form an ionic compound.

Carbon atoms can carry a charge as part of an ion, but only when they are part of a larger species (again, like $\ce{C2^{2-}}$) and they do not carry 4 additional electrons per atom. Check out cyanide as another good example of this type of polyatomic ion.

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    $\begingroup$ Just a tiny nitpick: The electron affinity of carbon is positive according to its definition, the energy required to detach an electron from the singly charged negative ion. This of course does not affect anything in your logic. $\endgroup$ – Martin - マーチン Oct 14 '14 at 7:30
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    $\begingroup$ Lead fluoride isn't ionic. $\endgroup$ – permeakra Feb 29 '16 at 9:20
  • $\begingroup$ @permeakra I'm not saying you're wrong, but can you provide a source for this? The only sources I can find say that it has "appreciable ionic character" unlike PbCl4, which is predominately covalent. $\endgroup$ – Jason Patterson Feb 29 '16 at 21:53
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    $\begingroup$ @JasonPatterson I guess, this belief comes from population analysis for PbF4 and PbCl4 that gives Pb charge +3.2 and +1.9 respectively. This analysis should not be trusted, since similar analysis for CH4 and CLi4 may give -0.47 and -0.50 respectively for carbon; and +0.2 for Li in LiH; which is an obvious BS. This comes from the flaws in calculation schemes for atomic charges. $\endgroup$ – permeakra Mar 1 '16 at 6:01
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Size decreases as one goes from left to right in perodic table. that makes carbon pretty small. So it will naturally be hard to take out electrons from it.

But what about accepting electrons but it does not do that either. the ansswer is "Fajan's rule".

It helps a bit to distinguish between covalent bonds and ionic bonds. but I would suggest that if you are in a small standard it might be a problem to understand Fajan's rule. Everybody has the same problem before Fajan's rule and Fajan's rule makes it a little easier to understand chemical bonding.

So for now just accept grow your knowledge about bonding understand Fajan's rule if possible.

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    $\begingroup$ It's Fajans' rules. You might want to tidy-up your answer: as it reads now, it's not easy to read/comprehend. $\endgroup$ – Todd Minehardt Aug 27 '16 at 18:56
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    $\begingroup$ This does not clearly address the original question. $\endgroup$ – Jon Custer Aug 27 '16 at 20:11
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    $\begingroup$ 60% of this answer is commentary on how awesome Fajan's rule is without even explaining it. Remove that, and you're left with a bunch of vague attempts at answering. Borderline NAA. If y'all flag I don't complain. $\endgroup$ – M.A.R. Aug 27 '16 at 22:20

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