Back in the day, alchemists apparently spent all their time attempting to turn base metals into gold. I say apparently, since that is the common perception of alchemists. In reality, most alchemists were devoted to what was often referred to as the “Great Work.” According to alchemical theories, much of what is created by nature is imperfect, and by applying the art of alchemy, natural substances can be brought to a higher state of perfection. Since gold was the height of perfection for metals, it was thought that base metals could be transmuted into gold merely by removing these imperfections. The Philosophers Stone wasn’t just about making gold or the “Elixer of Life”, it was a process/device/concept for bringing about perfection.
Color was very important to the alchemists, perhaps since it was one of the few clues with which they had to go on during their experiments. Colors and color changes were rigorously recorded and eventually incorporated into many of the alchemical theories. Certain color sequences were expected during the path towards perfection. In the transmutation of base metals into gold, for example, a color sequence of black to white to yellow to purple was thought to be required. Later theories redefined the sequence as black to white to red. While this fascination with colors may have led them astray on occasion, it has also led to a vast array of wonderful, full color illustrations, intricately drawn and full of alchemical symbolism. Here is one such example.
An experiment often touted as reminiscent of alchemy and demonstrating the relative ease of generating color changes involves the apparent conversion of a copper penny into silver or gold. Simply dissolve some Zn powder in warm NaOH solution, toss in a penny, and watch the penny turn silver as a coating of Zn forms on its surface. Heating the penny turns it gold as the Zn and Cu alloy to form brass. (My kids were duly impressed when I performed this experiment for them, informing me that it was “awesome”, but considering they said the same thing when I turned phenolphthalein red, the “awesomeness” bar may be set pretty low.)
But why does this work? I’ve coated pennies with silver and mercury before, but zinc is more electropositive than copper. Based on electrochemical potentials, zinc shouldn’t plate out on copper -- copper should be plating out on zinc. Apparently I’m not the only person to wonder about this, as web pages devoted to this effect can be found throughout the Internet, for example, here and here.
First of all, Zn dissolves in NaOH to make zincate ion and hydrogen.
Zn + 2OH- + 2H2O ----> Zn(OH)42- + H2
Okay, that’s straightforward chemistry. But how does the copper reduce the Zn(OH)42-? The answer is that it doesn’t. Surprisingly, copper is not the reductant . Copper is not oxidized and it does not go into solution. Copper’s role is to create a galvanic cell with the undissolved portion of the zinc powder. Zinc does not plate out on the penny until the copper is in direct contact with zinc metal. It turns out that zinc metal is the reductant, which sounds bizarre, at least to me. Once the zinc and copper metals are in contact and the galvanic cell is created, the zinc begins to oxidize…
Zn + 4OH- ----> Zn(OH)42- + 2e-
And its electrons are now available to reduce the zincate ion near the surface of the penny...
Zn(OH)42- + 2e- ---> Zn + 4OH-
So the overall reaction is:
Zn + Zn(OH)42- ----> Zn(OH)42- + Zn
Basically there is no net reaction! Which means the zinc is essentially just migrating from the surface of the zinc powder to the surface of the penny. At first glance, this would appear to violate the laws of thermodynamics, but obviously it doesn’t. I assume there is some sort of entropy effect here, perhaps related to the galvanic cell, but I don’t know the exact cause. In any case, this reaction appears just as mystical as alchemy itself!
If anyone happens to know any more about this process, I’d be happy to hear from you.
Showing posts with label transition metals. Show all posts
Showing posts with label transition metals. Show all posts
Tuesday, March 17, 2009
Friday, March 6, 2009
Colorful Chemistry
I suspect that one of the reasons I chose Inorganic chemistry as my major was due to the colorful chemistry of transition metal complexes. Organic compounds, at least the ones I saw during my first several years in college, were almost always white. And in those instances where some color was present, usually a rather boring pale yellow or brown, it was often due to the presence of impurities.
Unfortunately, the electronic transitions responsible for most of the transition metal colors are d-d transitions, which are generally forbidden under the rules of quantum mechanics, so it often requires fairly concentrated solutions to generate rich colors. There are the occasional exceptions – e.g. MnO4-, whose color is due to a quantum mechanically allowed electronic charge transition (the electron jumps from a metal orbital to an oxygen orbital) -- but generally, the extinction coefficients of most inorganic molecules are low.
So it’s rather ironic that a majority of the most deeply colored compounds are organic molecules. No forbidden electronic transitions here, just conjugated systems that can be tailored to absorb just about any wavelength of light in the visible and UV spectrum. This property has led to their use as dyes for over 4000 years. A list of the early dyes would include:
Alizarin – produced by the madder root
Carmine - obtained from the bodies of cochineal insects
Indigo – obtained from the indigo plant
Tyrian Purple – a brominated version of indigo, obtained from the Murex (a type of shellfish) in minute amounts, so quite expensive. Only affordable to the uber-wealthy, it eventually became seen as a symbol of royalty (thus the saying “born to the purple”). In Roman times, it was a capital offense to wear it if you were not a noble. Exposure of the dye to alkali turns it crimson, producing the color worn by Cardinals in the Catholic church.
Starting in the 1800s, many of these natural dyes were replaced with aniline based compounds produced from coal tar. Some of these are considered safe enough to eat, which is why Blue No. 2 (indigotine) is found both in your blue jeans and in your blue M&Ms. Not all of today’s dyes are synthetic. Carmine, which is still obtained from insects, is still used to impart a reddish color to some foods in the US, which is creating a bit of an uproar. Not surprisingly, the food industry is not enthused about telling consumers that some of their products are made from bugs.
The blue food coloring referred to as Brilliant Blue FCF (or Blue No. 1) has a noticeable side effect of which most parents are aware. Green poop. I remember the first time my six month old son presented me with such a gift. Unfortunately, he was suffering from some unknown intestinal disorder at the time which already had us a little worried. The only reason I didn’t immediately panic was that the bright Kelly green color was so artificial looking that it was hard to believe it was physiological in origin. Apparently purple goldfish crackers have Blue No. 1 in them.
Unfortunately, the electronic transitions responsible for most of the transition metal colors are d-d transitions, which are generally forbidden under the rules of quantum mechanics, so it often requires fairly concentrated solutions to generate rich colors. There are the occasional exceptions – e.g. MnO4-, whose color is due to a quantum mechanically allowed electronic charge transition (the electron jumps from a metal orbital to an oxygen orbital) -- but generally, the extinction coefficients of most inorganic molecules are low.
So it’s rather ironic that a majority of the most deeply colored compounds are organic molecules. No forbidden electronic transitions here, just conjugated systems that can be tailored to absorb just about any wavelength of light in the visible and UV spectrum. This property has led to their use as dyes for over 4000 years. A list of the early dyes would include:
Alizarin – produced by the madder root
Carmine - obtained from the bodies of cochineal insects
Indigo – obtained from the indigo plant
Tyrian Purple – a brominated version of indigo, obtained from the Murex (a type of shellfish) in minute amounts, so quite expensive. Only affordable to the uber-wealthy, it eventually became seen as a symbol of royalty (thus the saying “born to the purple”). In Roman times, it was a capital offense to wear it if you were not a noble. Exposure of the dye to alkali turns it crimson, producing the color worn by Cardinals in the Catholic church.
Starting in the 1800s, many of these natural dyes were replaced with aniline based compounds produced from coal tar. Some of these are considered safe enough to eat, which is why Blue No. 2 (indigotine) is found both in your blue jeans and in your blue M&Ms. Not all of today’s dyes are synthetic. Carmine, which is still obtained from insects, is still used to impart a reddish color to some foods in the US, which is creating a bit of an uproar. Not surprisingly, the food industry is not enthused about telling consumers that some of their products are made from bugs.
The blue food coloring referred to as Brilliant Blue FCF (or Blue No. 1) has a noticeable side effect of which most parents are aware. Green poop. I remember the first time my six month old son presented me with such a gift. Unfortunately, he was suffering from some unknown intestinal disorder at the time which already had us a little worried. The only reason I didn’t immediately panic was that the bright Kelly green color was so artificial looking that it was hard to believe it was physiological in origin. Apparently purple goldfish crackers have Blue No. 1 in them.
Thursday, August 21, 2008
Hooray! Aqueous Transition Metal Chemistry Talks
I really miss going to ACS meetings, but with chemistry not being as big a focus at my company anymore, it can be difficult to justify (to upper management) the expense of going to one unless there is a session directly related to my current project. I attended a session on reforming catalysts several years ago, but that’s been it for a while. Besides, my true love is aqueous transition metal chemistry and those types of talks can be few and far between. So imagine my surprise and delight when I ran across blogs discussing, not one, but two such talks being presented at this week’s national ACS meeting.
The first talk was mentioned on the Chemistry Blog, concerning a study to use chelating agents like EDDA to extract actinides from nuclear waste. YES! Good, old-fashioned aqueous transition metal chemistry. Okay, actinides aren't strictly transition metals, but they're close enough. I did find it amusing when Mitch wrote "These types of systems look promising and are a nice upgrade over traditional old-school extractors like tributylphosphate(TBP)." Any reference to classic chelating agents as "new-school" is music to my ears. To be honest though, I would have thought that these chelators, especially EDDA, would have been tested for this application a long time ago. At least it shows there are still applications for this type of chemistry yet to be discovered.
The second talk involved the use of Fe-TAML complexes to facilitate the oxidation of organic pollutants by hydrogen peroxide. Apparently these complexes are good mimcs of peroxidase enzymes. Here is the structure of the compound taken from the Collins group website:
Although I haven't had the chance to read the mechanistic and kinetic work on these reactions yet, this seems like a classic example of the TAML group basically locking the Fe into a square planar configuration and keeping it stable while it’s shuttling back and forth between the +2 and +3 states as it facilitates the transfer of electrons between the H2O2 and the target molecule. I've always had a soft spot for chelating agents that lock themselves around metals in tight, specific configurations, significantly changing the metal's properties in the process. My first undergraduate project used an amine oxime to wrap itself around cobalt in a manner resembling the structure of cobalamin (B12 vitamin). Classic aqueous transition metal chemistry. It brings a tear to my eye.
The first talk was mentioned on the Chemistry Blog, concerning a study to use chelating agents like EDDA to extract actinides from nuclear waste. YES! Good, old-fashioned aqueous transition metal chemistry. Okay, actinides aren't strictly transition metals, but they're close enough. I did find it amusing when Mitch wrote "These types of systems look promising and are a nice upgrade over traditional old-school extractors like tributylphosphate(TBP)." Any reference to classic chelating agents as "new-school" is music to my ears. To be honest though, I would have thought that these chelators, especially EDDA, would have been tested for this application a long time ago. At least it shows there are still applications for this type of chemistry yet to be discovered.
The second talk involved the use of Fe-TAML complexes to facilitate the oxidation of organic pollutants by hydrogen peroxide. Apparently these complexes are good mimcs of peroxidase enzymes. Here is the structure of the compound taken from the Collins group website:
Although I haven't had the chance to read the mechanistic and kinetic work on these reactions yet, this seems like a classic example of the TAML group basically locking the Fe into a square planar configuration and keeping it stable while it’s shuttling back and forth between the +2 and +3 states as it facilitates the transfer of electrons between the H2O2 and the target molecule. I've always had a soft spot for chelating agents that lock themselves around metals in tight, specific configurations, significantly changing the metal's properties in the process. My first undergraduate project used an amine oxime to wrap itself around cobalt in a manner resembling the structure of cobalamin (B12 vitamin). Classic aqueous transition metal chemistry. It brings a tear to my eye.
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