PDA

View Full Version : Alfred Werner and the Semi-factual Story of Coordination Chemistry



jubei33
25th April 09, 06:58 PM
In 1913, a Swiss chemist named Alfred Werner (http://en.wikipedia.org/wiki/Alfred_Werner) was awarded the Noble Prize in Chemistry for his work on what would be called coordination chemistry (http://en.wikipedia.org/wiki/Coordination_chemistry), which would lead to a new understanding of how chemicals bond together. Coordination theory describes the nature of bonding in transitional metals and the formation of complexes, which at the time seemed to follow bizarre and unpredictable patterns. Atoms in groups 1-7A followed somewhat predictable patterns in their bonding as shown by classical experiments. For example, atoms in group I took on +1 charges and bonded once to other more negatively charged atoms like the group 7A halogens, but it remained a mystery how the transition metals bonded and why they had so many oxidation states (http://en.wikipedia.org/wiki/Oxidation_states). The shape of salts like (CoCl3 * 6 NH3) was still undetermined and throughout much of the 1800’s one popular theory emerged: chain theory. It was supported by some of the most powerful chemist-sorcerers at the time, including Werner’s chief rival: S.M. Jorgensen (http://en.wikipedia.org/wiki/Sophus_Mads_J%C3%B8rgensen).

http://i131.photobucket.com/albums/p297/jubei33/periodic_table.gif

Jorgensen believed that the ligands in compounds like (CoCl3 * 4 NH3) were arranged in chains, that is, bonded to each other in some fashion. The main point being that the atoms would follow known valence rules at the time, especially Kekule (http://en.wikipedia.org/wiki/Friedrich_August_Kekul%C3%A9_von_Stradonitz)’s principle, which abstracted the number of times a compound could bond from known chemical reactions. Though useful, it ran into problems when trying to describe why atoms with larger electronic configurations bonded in so many different arrangements. Transitional metals in particular confounded these rule sets.

http://i131.photobucket.com/albums/p297/jubei33/chaintheorypossibilites.jpg
Possible chain theory configurations of (CoCl3 * 4 NH3).

http://i131.photobucket.com/albums/p297/jubei33/vanadiumrainbow.jpg
A vanadium rainbow: Potassium permanganate is added to a Vanadium II solution. Over time it separates into different states: V +2 is violet, V +3 is green, VO +2 is blue, VO2 +1 is pale yellow, MnO2 is brown and MnO4 –1 is pink.

Werner, however, proposed a different theory that relied on the concept that cobalt (in the above compound) could have more than three bonds and that the ligands would be centered around cobalt in an octahedrally shaped crystal. According to his theory, a compound like the above due to its structure would have two possible conformations: a cis isomer (http://en.wikipedia.org/wiki/Cis-trans_isomerism) (with chlorine atoms on adjacent vertices) and a more stability favored trans isomer (with the chlorine atoms on opposite side of one another). Interestingly, the two are identifiable by their color with the trans compound being green and cis being a delightful purple color.

http://i131.photobucket.com/albums/p297/jubei33/greentranspurplecis.jpg

Naturally, this caused controversy amongst chemists and the debate began. At the time only the structurally favored green trans compound had been synthesized, leaving the more difficult cis compound to be thought of as non-existant. Whenever Werner published results that seemingly confirmed his theory, Jorgensen was there to propose a counter theory in favor of the more popular chain theory. Eventually, Werner was able to prove his case conclusively through a variety of methods like optical resolution of the compounds and electrical conductivity measurements. The capstone, as the story goes, was his synthesization of the elusive purple cis isomer of [Co (NH3)4 Cl3] and sending a sample through the mail to Jorgensen.

http://i131.photobucket.com/albums/p297/jubei33/makepurplecis.jpg
Werner used this clever method to synthesize his purple cis isomer. By adding HCL at 0C, carbon dioxide is released and chlorine atoms in solution replace the oxygen atoms lost. (Note: picture does not show positive charge on Cobalt atom)


End part I

Ajamil
26th April 09, 09:27 AM
Science needs more trashtalking as they work on their opposing theories. Otherwise great write-up.

Cullion
26th April 09, 10:16 AM
That raises many, many basic questions about Chemistry as I haven't studied it since I was 18.

Here are a couple:

Why does is trans-isomer in the above example more stable than the cis isomer?

The two isomers are different colours, what causes this and how easy is to predict what colour a compound will be from knowing it's molecular structure ?

RaiNnyX4
26th April 09, 03:52 PM
That raises many, many basic questions about Chemistry as I haven't studied it since I was 18.

Here are a couple:

Why does is trans-isomer in the above example more stable than the cis isomer?

The two isomers are different colours, what causes this and how easy is to predict what colour a compound will be from knowing it's molecular structure ?

I'm a Chemist too. But whereas Jubei's specialty seems to be Organic/Inorganic in nature, my studies were focused on the Math/Physics side of Chemistry.

I don't remember too much from my Inorganic Chemistry days but my hypothesis for the greater stability of the trans isomer would be due to Chlorine's high electronegativity. Usually, Chlorine will have an apparent negative charge when bonded with other elements. Remembering that like charges repel, it's natural for a compound to seek a conformation where similar charges will be separated by as much distance as possible. In the cis conformation, the two Chlorine atoms are separated by 90 degrees, whereas in the trans they are 180 degrees from each other.

As for the second question, I'm sure I learned the answer to that at some time but I honestly don't remember.

jubei33
26th April 09, 04:07 PM
all of the Halogens in group 7a are electronegative atoms, which try their damned hardest to get that last electron to fill their outer electron orbital (octet rule (http://en.wikipedia.org/wiki/Octet_rule)). They're desperate and they'll do anything to get what they want, even if they have to steal from one of their own. Fluorine is considered the most electronegative, but don't count chlorine down, he'll gut you too if you're not careful.

More seriously, Trans compounds like the above show greater stability because the electronegative atoms are furthest away from each other, leading to a greater ease of formation and lower activation energy comparatively.

Now the colors in compounds come from their electronic structure as well. When an atom is blasted by some beam of energy (in this particular instance its usually a photon), electrons get excited. To relieve this extra energy, electrons are shifted from higher to lower orbitals and the excess is released as photons. Oftentimes this happens at predictable wavelengths, which consequently result in wonderful colors. This was one of the big questions I had as a kid too, I loved seeing the colors change in chemistry, it was bizarre and magical.

edit: rainy's got the right as well. posted while I was posting..


That raises many, many basic questions about Chemistry as I haven't studied it since I was 18.
Ask away, I doubt I could answer all of them, but all we can do is try.



Science needs more trashtalking as they work on their opposing theories. Otherwise great write-up.
then you'll like part two, I think. It ends in a manly embrace.

Cullion
26th April 09, 04:15 PM
More seriously, Trans compounds like the above show greater stability because the electronegative atoms are furthest away from each other, leading to a greater ease of formation and lower activation energy comparatively.

Ok, I was taught this once and now you've reminded me, thanks.



now the colors in compounds come from the electronic structure as well. When an atom is blasted by some beam of energy (in particular a photon), electrons get excited. To relieve this extra energy, electrons are shifted from higher to lower orbitals and the excess is released as photons.

Why do we get different colours for different Isomers which are comprised of the same atoms? Presumably the shape of the molecule affects which orbitals electrons will shift between when a photon hits them. Is predicting this mathematically a very complex problem of quantum mechanics ?

Do we reach a point in molecule complexity where the maths is just not solvable and it all has to be determined by experiment ?

jubei33
26th April 09, 04:28 PM
usually the electron transitions are d-d transitions. The big thing werner and his chain theory buds were missing was a good theory of the atom, which didn't come until later. D-orbitals are the 'sandwich meat' of complex compounds and transition metals. Even though they're the same, different locations of ligands means slightly different electronic structure. To predict it and see the outcome we can use spectroscopic methods like UV-Vis and tanabe-sugano charts (http://en.wikipedia.org/wiki/Tanabe-Sugano_diagram).


Do we reach a point in molecule complexity where the maths is just not solvable and it all has to be determined by experiment ? possibly, I don't know.

edit: In these compounds, electrons can also be transferred to a ligand's orbitals as well.

Ajamil
27th April 09, 05:58 PM
Ever hear of a book called Weighing the Soul? It detailed several high profile cases of science reaching a split in theory, and the experiments/arguments scientists for each side went through.

SFGOON
27th April 09, 11:41 PM
So, don't chemicals piss you off when you iodine-halogenate an irritant to turn it into an antagonist for the VR-1 pathway, not realizing that a carbon/iodine bond has a heat of formation of +15, so when you take it out of the fridge and dissolve it in glycerine all the iodine (the fat lazy excuse for a halogen) wanders off and turns your 6-iodo-caspicium into plain ol' caspicium.

Then you spray it in your eyes and hurt yourself?

jubei33
29th April 09, 03:52 PM
So, don't chemicals piss you off when you iodine-halogenate an irritant to turn it into an antagonist for the VR-1 pathway, not realizing that a carbon/iodine bond has a heat of formation of +15, so when you take it out of the fridge and dissolve it in glycerine all the iodine (the fat lazy excuse for a halogen) wanders off and turns your 6-iodo-caspicium into plain ol' caspicium.

Then you spray it in your eyes and hurt yourself?
Hey, you OK dude? Don't blind yourself.


Ever hear of a book called Weighing the Soul? It detailed several high profile cases of science reaching a split in theory, and the experiments/arguments scientists for each side went through.

no, never read it. Why don't you write a synopsis/review.