crazy ox-states

[akcapr]

my teacher claims no one elements can gain/lose more than 4 electrons- how is this possible since oxidation states get higher like 6? where am i confusing myself and wats the jist?

[woelen]

Oxidation state is just a book-keeping device. For example in the compound CrO3, the metal chromium is in the +6 oxidation state, but this does not mean that it gas lost 6 electrons. The bonds in CrO3 are mostly covalent, with just a small charge-bias towards the oxygen atoms.

 

Only for purely simple ionic compounds, the oxidation state equals the charge.

 

E.g. in CaO, the calcium really lost two electrons and exists as Ca(2+) ion and the oxygen gained two electrons and exists as O(2-) ion. In a compound like HgO, things are quite different. The bond in this compound is not fully ionic, but something between purely covalent and purely ionic.

 

So, comparing these two compounds:

 

(2+)Ca O(2-)

(+)Hg....O(-)

 

In CaO, the ions are completely disconnected, in HgO, there is still some bond between the Hg-atom and O-atom, but charge is distributed with a positive end at the Hg-atom and a negative end at the O-atom. However, the Hg-atom does not have a +2 charge.

 

This effect becomes stronger with increasing oxidation state. So, for a compound like CrO3, the Cr-atom carries a slight positive charge and the O-atoms carry a slight negative charge.

 

Why is the concept of oxidation state introduced? It provides a good way of bookkeeping the level of oxidation (or reduction) of an element and it allows one to see whether an element is in a common oxidation state or a very special one. With some experience, you can see at once, whether a compound is as expected or something very special.

[akcapr]

so in the chromium then, in the +6 oxidation state, how did it come to that oxidation state without losing so many electrons

[woelen]

so in the chromium then, in the +6 oxidation state, how did it come to that oxidation state without losing so many electrons

As I wrote before, it just is a matter of bookkeeping and does not reflect the real structure of this covalent compound.

 

There are some bookkeeping rules and for this discussion I repeat the one, which is important here:

Oxygen has oxidation state -2 in almost all of its compounds. Exceptions are peroxo-compounds, superoxo-compounds, ozonide-compounds.

 

Now back to CrO3. This is a neutral species. Oxygen has oxidation state -2, there are three oxygens, so chromium has oxidation state +6 in order to obtain neutrality.

 

Another example: NO2(-) ion, nitrite. Oxygen has oxidation state -2, there are 2 of these. The total charge equals -1 and there is one N-atom. Make it +3 and you end up with the charge equal to -1. So, N in nitrite has oxidation state +3. The N in the brown nitrogen dioxide (NO2 without charge) has oxidation state +4. The N in the nitronium ion (NO2(+)) has oxidation state +5.

 

Things become even more interesting, such as in a salt like nitronium perchlorate (NO2 ClO4, with the NO2(+) ion and the ClO4(-) ion). I'll leave it as an exercise for you to compute the oxidation state of the chlorine atom in this salt. The actual charge on that atom, however, is much lower, because it is covalently bonded to the oxygens.

[akcapr]

ok thx, but how does chromium balance out the +6 of the 3 oxygens if it itself isnt literally loosing somany electrons?

[woelen]

ok thx, but how does chromium balance out the +6 of the 3 oxygens if it itself isnt literally loosing somany electrons?

Well, I have the idea that you still do not really grasp the concept.

 

Another time, oxidation state is just bookkeeping, it does not reflect real structure.

 

I'll try to explain with two very simple compounds:

 

CO (carbon monoxide)

CaO (calcium oxide)

 

The first is a purely covalent compound. This means that the C atom and O atom are bound tightly and that both atoms share electrons. Two electrons from the O and two electrons from the C are shared between the C and O atoms. Of each of these 4 electrons you cannot say to which atom they belong, all 4 of them belong to both atoms.

 

Now, if you look at CaO, this is a purely ionic compound. It consists of Ca(2+) ions and O(2-) ions, packed in a crystal lattice, but nevertheless, these ions exist as real entities.

 

In the last example, oxidation state really matches the real physical structure. In the first example, there is no ionic thing at all.

 

However, in both compounds, the oxidation state of the oxygen atom is said to be -2. It is just a matter of appointment to say this. Both compounds are neutral species, so for the other element the oxidation state must be +2. Again, for CaO this is the real charge on the calcium ion, but for CO again, it the oxidation state of carbon is said to be +2.

 

Now back to CrO3. There is no actual +6 charge on the chromium, we simple say it has oxidation state +6, because we say that oxygen has oxidation state -2 in its compounds (except the few I mentioned earlier).

 

Now, the link with real chemistry is just that we order elements from most electronegative to most electropositive. Electronegative elements usually have negative oxidation states. Fluorine is the most electronegative element and it is said to have oxidation state -1 always, except in F2. Oxygen has oxidation state -2, except in the compounds I mentioned earlier and in OF2, where it has oxidation state +2 (this must be the case, because of the rule for fluorine). Chlorine is the next electronegative elements and it has oxidation state -1 usually, but in its compounds and ions with fluorine and oxygen it has a positive oxidation state. The oxidation state, hence, tells something about how much an element is oxidized. The higher the number, the stronger it is oxidized. So, Cr2O3 contains a much less oxidized form of chromium than CrO3. When you know the rules and have some experience, you can see at once, whether a compound is a strongly oxidized compound or not.

 

You understand? Just plain arithmetic.

[jdurg]

To put it simply, oxidation state and charge are two completely different things. Sometimes they happen to match, but they are NOT the same thing.

[latentheat]

You could basically say that ox state just tells you how many bonds have been made, and only reflects lost or gained electrons in ionic bonds.

[akcapr]

Things become even more interesting' date=' such as in a salt like nitronium perchlorate (NO2 ClO4, with the NO2(+) ion and the ClO4(-) ion). I'll leave it as an exercise for you to compute the oxidation state of the chlorine atom in this salt. The actual charge on that atom, however, is much lower, because it is covalently bonded to the oxygens.[/quote']

 

 

im sure im wrong but..8+? my reasoning is each oxygen makes 2 bonds ad theres 4, so.. 8

[woelen]

im sure im wrong but..8+? my reasoning is each oxygen makes 2 bonds ad theres 4, so.. 8

In fact, it is +7. ClO4(-) has a charge equal to -1. You have four oxygens, each having an oxidation number equal to -2, that makes up a total of -8. In order to have a net charge of the ion, the chlorine must be +7.

 

Remember, oxidation state and ionic charge at an atom-basis are two completely different things, but the following is true: The sum of all oxidation numbers of all atoms in an ion or a molecule must be equal to the total charge of the ion or molecule.