Yes, I know; reading this statement might make people think that megagallons of ethanol are coursing through my arteries. It sounds quite counterintuitive. But I have become convinced that this is the case. I have not explored the matter in exhaustive detail, and some questions still linger. The explanation can get long-winded and I will reserve it for sometime when my own thoughts are clearer. But here's the summary.
There is a very simple reason why all transition metals will prefer only the +1 state in their complexes; Pauling's electroneutrality principle.
More simply put, consider this; won't it be improbable for a charge of +2 or +3 to accumulate on a metal ion? There will be repulsive destabilization. We know that charge tries to disperse itself and stabilize any system. So this positive charge will try to redistribute itself and get delocalized, thereby making the system close to electrostatically neutral. Where can it do this? Well, we don't want to forget 'the other half'- the ligands. What this means is that it's actually the ligands that are getting oxidised by helping to delocalize the positive charge. So even in a compound like AuCl3, the Au has a +1 charge, and it's really the chlorides that are getting oxidised.
The whole root of the problem is that chemists' definition of oxidation state actually is quite arbitrary, based on qualitative notions of electronegativity. We think that since chlorine is more electronegative than gold, it will carry all the negative charge. But Pauling's electroneutrality principle tells us otherwise.
However, this is not something we usually think about. Consider the simple ammonium cation. Organic chemists depict it with the positive charge on nitrogen. But it's a well-known fact that the positive charge in fact is distributed over the hydrogens. Any system will try to stabilize itself by dispersing charge as much as it can. That's what the ammonium cation does, nicely dispersing the positive charge over the hydrogens.
We might say; but how can halogens get oxidised? They are supposed to be among the most strongly oxidising elements. But note, that this statement applies to diatomic halogen molecules, not to halogen ions in solution.
You can take my word that even the fiercely electronegative CF3- anion in the complex [Cu(CF3)4]- is oxidised, and that the copper exists as Cu(I) in this state. Even a look at the formula suggests otherwise; the Cu must be existing as Cu(III). But you don't have to take my word for it. Consider this publication in Angewandte Chemie. There was also a small debate, published in another issue of the same journal, which further combs some fine points and definitions and clarifies things. High level calculations strongly suggest that the Cu exists as Cu(I).
High level calculations, but more logically, the electroneutrality principle and basic concepts of electrostatics.
The point is, many things in chemistry are defined either arbitrarily, or with respect to some standard. Based on the existence of another chemical species, many things can change for a given species. Given the right conditions, almost anything can get oxidised or reduced. Given the right conditions, anything can be made stable. After all, nobody had thought that stable carbocations could exist. But George Olah opened up an entire new field of carbocation chemistry with his pioneering work. In chemistry, there are no absolutes. Which again proves why it is so interesting.
Actually, it's simple to think about. Usually we assume that the metal is changing its oxidation state. Instead of that, let's adopt the other perspective; the metal makes the ligand change its oxidation state. The result is the same. We still see AuCl3 in solution. But it's really the Cl- that has changed, and not the Au.
Some questions still linger; then why do all these metals seem to exist in different oxidation states. More importantly, how can one prove this contention by experiment? After all, ours is an experimental science.
But even in the absence of definite experiments, if one ad hoc concept is taken for granted, then another ad hoc concept surely cannot be worse. Especially if it sounds more logical and convincing. Like all metals existing only in +1 oxidation state (or something close to that).
Important addendum: I should mention that when I say +1, the actual number is better defined as "close to +1". Calculations show that it can change from +0.7 to about +1.2 (this bears on the question d_orbital asked in the comments) depending on the ligand environment. Thus, in hemoglobin for example, in the oxidised and reduced form, Fe won't exist as exactly +1 both of the times. Its oxidation state will fluctuate between decimal places close to one.
And, as Mitch quite rightly says, in most of our daily life, metals have an unequivocal oxidation state of 0!!
Field of Science
-
-
-
Political pollsters are pretending they know what's happening. They don't.4 weeks ago in Genomics, Medicine, and Pseudoscience
-
-
Course Corrections5 months ago in Angry by Choice
12 comments:
Markup Key:
- <b>bold</b> = bold
- <i>italic</i> = italic
- <a href="http://www.fieldofscience.com/">FoS</a> = FoS
Subscribe to:
Post Comments (Atom)
Interesting proposal. Several questions/comments come to mind.
ReplyDeleteIt was my understanding that formalism was the act of taking a compound and breaking it into its ionic components and assigning Lewis charges as such. As I came to know it, this concept wasn't to emulate Nature but to act as a symbolism for chemists to communicate with one another. Sort of like drawing an amide bond with the double bond between the carbon and the oxygen, even though we all know it is shared, albeit unequally, between all three atoms. Thus it is a way for us to realize the electronic differences between CuL and CuL3
This explanation seems reasonable enough when talking about CuLn3 where all Ln = each other. How does one account for heteroleptic complexes where the ligands are all different. Does one ligand exists as +0.9 and another at +1.1? Or does the copper exists as +1.2, etc? Since some ligands are more electronegative than others wouldn't we expect an inequality in the degree that the ligand is oxidized (by your explanation)?
How does this explain tetrafluoroborate type metal complexes where the metal is said to be in the (+) state [ie Rh(+)]. In these complexes is the metal really acting on the borate in an intimate contact pair?
Finally how does this account for ionic solutions? If you dissolve MgBr2 in a ionic solution, does Mg++ not exist? Instead are we left with Mg(0) and 2Br+?
In my day to day life all the metals I interact with have an oxidation state of zero, so your claim seems very-very far fetched to me.
ReplyDeleteMitch
d_orbital: I will respond to your questions in detail later. I have made a point in the addendum for now.
ReplyDeleteMitch: :) Absolutely true!
Thanks both of you for reminding me about these points; I updated the post.
One immediate concern I have with this is that your argument that it seems improbable that a positive charge can build up on a metal center due to repulsive interactions.
ReplyDeleteSince all the positive charge in the metal is localized in the nucleus and the oxidation state is merely a measure of lack of electrons, how does the repulsive interactions increase as charge increases?
I'd like to quickly expand on my first paragraph in my initial post:
ReplyDeleteIn Hegedus' excellent book "Transition Metals in the Synthesis of Complex Organic Molecules" he states in chapter 1, "...These formalisms are just that – formalisms – not reality, not the "truth", and in some cases, not even chemically reasonable. However, by placing the entire organic chemistry of transition metals within a single formalistic framework, an enormous amount of disparate chemistry can be systematized and organized, and a more nearly coherent view of the field is available."
In the very next paragraph, he immediately launches into a discussion of the concept of the oxidation state. He goes on to say, "The oxidation state of a metal is defined as the charge left on the metal atom after all ligands have been removed in their normal, closed-shell, configuration – that is, with their electron pairs. The oxidation state is not a physical property of the metal, and it cannot be measured. It is a formalism which helps us in counting electrons, but no more."
Whew!
Stewiee: I am not sure I understand this...one can also make a similar argument for organic reactive species, such as carbocations. The second point that you made is quite pertinent; it's a lack of electrons that creates an electron deficient species. Ligands around such a species will try to stabilize it. I personally don't know of any tri or dipositive organic cation where the charge )or electron deficient character if you will) is localized.
ReplyDeleteOne very important point I forgot to mention in the post is that the whole concept of oxidation state that I have seen defined in many textbooks rests on the bonding being 100% ionic. That itself should raise a red flag, because 100% ionic bonding almost never exists. I believe we can accept the argument even if we don't go any further, but just think about the problem with purely ionic bonding.
d_orbital: That's a coincidence! As far as I know too, Hegedus's book is the only organometallic book which makes this matter clear. Thanks for the reference.
I don't know much about ionic solutions, but let's say we have MgBr2 in water. There, both the "Mg2+" and the 2Br- ions are going to be solvated...in which case the ligands are water and they similarly are oxidised.
To be honest, I don't know about the Rh(+) ion and what you say may certainly be the case.
In case of different ligands, yes, they would be differentially oxidised.
Great Post!
ReplyDeleteThe oxidation state and valency formalisms(i.e. they are just formalisms and they don't work in organometallics) were the main ideas throughought "The Chemistry of elementoorganic Compounds" and " Metals in Biomolecular Chemistry" courses I took last semester.
I think that basic idea that you are trying to convey is interesting - but the way you are doing so is incorrect. The difference that you wish to describe is that the formal oxidation state of an atom is different from its effective oxidation state - or as is generally termed - the effective nuclear charge of the atom. The effective charge is essentially the charge that the atom 'seems' to possess even though it does not exist as a naked ion. This is, of course, not true for true ions in the gas phase but this is not what this is about. I think it is best to consider that the effect of reduction/oxidation is compressed by the ability to delocalize charge over the rest of the molecule (to a certain extent).
ReplyDeleteSo, for example, if we consider the possible formal oxidation states of iron, we would say that the charge changes in integer steps: 0, +1, +2, +3, +4. However, in reality, the metal does not carry an effective charge of +1, +2, +3, +4 - the changes between these would be far more compressed such that although the Fe(IV) species isn't really Fe(4+), it IS more positively charged than a comparable Fe(III) complex.
The biggest utility of defining a formal charge, though, is not in trying to accurate define the charge on the atom - it is more a way to represent the correct limiting electronic description for a particular ion. If we use the iron example again and we look at the difference between Fe(II) and Fe(III) complexes, there is definitely a larger charge at the metal in the Fe(III) complexes (as evidenced by a whole host of techniques) but the really big difference is that the Fe(II) complex is best described as having a 3d6 configuration whereas the Fe(III) has a 3d5 configuration.
One definite problem with your analysis is that it to take the principle of electroneutrality a bit out of context - in the sense that it doesn't actually mean that charge distribution has to be even. The real issue is that the chemical potential must be equalized throughout the molecule but that very well might involve having one atom being very positively charged with very little contribution from surrounding atoms - especially in transition metal complexes since charge delocalization (or redistribution) requires covalent interactions. In purely organic molecules, bonding is predominantly covalent in nature and therefore charge redistribution tends to be relatively facile. In transition metal complexes, electron delocalization can be very different from one system to another and therefore charge redistribution can differ substantially.
On a final note... theoretical chemists continue to have difficulty in accurately defining so-called atomic charges in covalent system and different approaches often yield wildly different answers. Recent advances in this area such as Bader's "atoms in molecules" approach have provided greater rigour in this area but there is still no single 'correct' way to assign charges when electron density is dispersed over the whole molecule (and necessarily between atoms).
I need to add one more thing to my previous comment since I've had a chance to read through the original post again and read my response.
ReplyDelete"...then why do all these metals seem to exist in different oxidation states. More importantly, how can one prove this contention by experiment?"
Okay, you've actually got lots of options here.
(1) If you are trying to determine effective charges, this can be done using X-ray photoelectron spectroscopy (XPS) and/or X-ray absorption spectroscopy (XAS) to look at the ionization energies of core orbitals, which are quite sensitive to the electron cloud.
(2) One can look at ligand field spectroscopy (in the UV/Vis/nIR region) to look at d-d transitions, from which one can derive d-orbital occupancies - i.e. you can tell the difference between a 3d5 and 3d6 ion quite easily in many cases (but not always!). You can use the equivalent transitions in XAS to do the same thing - and it's even a bit easier to do (the analysis) even though the data is harder to get.
(3) You can use electron paramagnetic resonance (EPR) to determine the amount of electron delocalization between atoms in cases where you have unpaired electrons and thus directly determine how covalent the bonding is.
(4) You can use valence photoelectron spectroscopy (PES) to look at the composition of bonding molecular orbitals and determine metal and ligand contributions to them.
(5) ...
I'll stop here - but there are definitely ways to address this - and many people deal with these issues in their work.
The upshot on about 30 years of work is the following: For the most part the redox chemistry of transition metal species indicates predominantly metal-based redox changes but in some cases, that isn't the case. I'll point you to the recent work of Karl Wieghart both on the identification and characterisation of Fe(V) and Fe(VI) complexes (examples of how you can actually figure out what formal oxidiation state is most appropriate from an experimental standpoint) as well as looking at a 30 year old question of metal-dithiolene complexes and whether their redox chemistry is best described as metal or ligand-based (original disagreement between Dick Holm and Harry Gray). Sorry no specific references but these should be quite easy to find - all within the last year or so.
Anon: Thanks very much for the extensive references! It would take a while to chew on all this info.
ReplyDeleteAbout the change in charge, yes, there is some change, but the perceived change of "one unit" is too large. Also, even though the difference in the d shell count is true, the real question is of how to distribute the electrons. Of course, in one case we have a closed shell system and in the other we have an open shell system (and this can naturally be experimentally borne out because of the difference of the unpaired electron), but it still leaves the question of how we define the charge and as you said, there is no ultimate way to do that.
Tynchtyk: Thanks
Your discussion on a very relevant topic seems to have died more than a year ago, but you may have been onto something...
ReplyDeleteCheck out the June 5 issue of Nature (vol. 453, News & Views on p. 735 and Letter on pp. 763-766).
Electroneutrality principle, introduced more or less one hundred years ago by I. Langmuir, that says that any atom composing a molecule try and tend to reduce its own eletrical charge and cannot remain with a charge higher than + or - 1, and the derived rule of adiacent charge adfirming that two adiacent atoms, connected by any type of bond at the usual atomic distance less than 5 A°, cannot bear a charge of the same sign and again with a value of less or equal to + or - 1, are strongly criticized from at least 15 years by a lot of physical-chemical scientists, the so-called Zintl-Klemm followers.
ReplyDeleteThese technicians every time try to demonstrate that mastodontic polimerized inorganic anions with an enormous global negative charge (e.i. from 5- to 15- for the negative ione and more than about 2- for each adiacent atom) are usual in the intermetallic and metal-oxide compounds and that they are stable against any electroneutrality myte or rule.
They consider that more important for the stability of an atom is to reach and fill completely electronic s-p orbitals with pairs of electrons assumed from everywhere (and relevant enormous negative charge assumed), and that Langmuir and Pauling principles are nothing or old fantastic stories.
These intermetallic and oxide compounds with enormous polymerized anions are proposed continuously and there is a race to invent any type of (probably inexistent or unstable) inorganic anions.
Now the situation from the chemical point of view is totally confused and unundestandable, and seems probable that we must leave again the "first priciples" or various versions of DFT, unable to clarify the structure of a chemical substance, an impossible task after this brain mixing, but to start and study again chemistry from the beginning for the beginners.
Thanks and kind regards.