Volume 77, Number 3
CENEAR 77 3 p.
As a child, I loved nothing so much as to see an object glowing in the dark. Back then, before such items were commonplace, I scoured toy stores in search of elusive phosphorescent balls, paints, or rubber bugs. I quickly learned to eschew the common "fluorescent" toys that required black lights--it was long-lasting, unaided glow I was after.
When I was a teenager, another wonderful glowing species began to appear at concerts and outdoor festivals: necklaces and ropes filled with mysterious liquids that shone brilliant greens, yellows, and blues. The glow lasted only a few hours, and unlike phosphorescent toys, once their light faded, they were irreversibly spent. My friends and I put them in the freezer, trying to extend their luminescence just a little longer.
And, of course, there were the light sticks, available at drug stores: short, plastic, tapered rods that contained a bizarre, brilliant green liquid. It was great fun to bend the plastic stick slightly, and hear the crunch of a small, liquid-filled glass capsule breaking within. As the two ingredients blended, an eerie bright green glow spread throughout the stick.
No longer a curiosity, nowadays chemiluminescent products are fabricated in every color and shape imaginable. But despite my fascination and chemiluminescence's ubiquity, I had never learned the chemistry that produced the glow and made the different colors. The shine of the green light stick and blue necklace was still largely a mystery.
Of course, as long as there have been fireflies glowing in the dark, humans have been fascinated by "cold" light. The firefly's glow mechanism, which hinges on the oxidation of firefly luciferin, is incredibly efficient--80 out of 100 reacting molecules go on to produce a photon of light.
In the early 1960s, when scientists took the first steps toward developing their own version of a firefly, they knew what was required: a molecule that radiates light when it's excited and an energy source to excite the molecule. There are numerous possible energy sources, such as light, heat, and electricity. In chemiluminescence, that source is a chemical reaction.
But the reaction has to generate a huge bundle of available energy, and transfer it as a package instantaneously to the fluorescent molecule without radiating any as heat. Known examples of such perfectly orchestrated reactions were few.
Also in the early 1960s, Edwin A. Chandross, a young chemist at Bell Labs in Murray Hill, N.J., was searching for a general way to explain chemiluminescence. Peroxides, with their potential to liberate large amounts of energy during some chemical reactions, seemed to be likely participants.
After a number of experiments, he found to his great excitement that oxalyl chloride mixed with hydrogen peroxide and a fluorescent dye produced chemical light. The efficiency was only about 0.1%, but it was the foundation from which sprang modern chemiluminescence. Chandross, unaware of the powerful potential of his discovery, never patented it.
At about the same time, chemist Michael M. Rauhut was manager of exploratory research at American Cyanamid in Stamford, Conn. He and his colleagues corresponded with Chandross about his oxalyl chloride chemistry, then went to work on the reaction--studying it and looking for avenues that would produce chemical light intense enough to be of practical use.
Rauhut and his colleague Laszlo J. Bollyky developed a series of oxalate esters. Ultimately, Rauhut designed a phenyl oxalate ester that, when mixed with hydrogen peroxide and a dye, gave a quantum yield of 5\--not as efficient as a firefly, but still brilliantly useful. They dubbed it Cyalume, and it became the trademark name for American Cyanamid's chemical light products.
"It was a great project," Rauhut, who is now retired, recalls fondly. "There were a lot of surprises."
The mechanism that he and other researchers have proposed for the process still stands as the best candidate: The oxalate ester and H2O2 react with the help of a salicylate catalyst to form a peroxyacid ester and phenol. The peroxyacid ester decomposes to form more phenol, and most important, a highly energetic intermediate, presumed to be a four-membered ring dimer of CO2. As the cyclic dimer decomposes into two CO2 molecules, it gives up its energy to a waiting dye molecule, which then fluoresces.
The group went searching for fluorescing dyes to make different colors. For example, the common green in most light sticks comes from 9,10-bis(phenylethynyl)anthracene, and 9,10-diphenylanthracene gives blue. "We invented a beautiful yellow," Rauhut remembers.
American Cyanamid eventually sold its chemical light division in 1993 to Springfield, Mass.-based Omniglow, a manufacturer of chemiluminescent products.
E. Earl Cranor, head of Omniglow's R&D, continues to develop new commercial uses for chemical light. His latest project is a light stick that works at below-freezing temperatures. He's also always seeking greater efficiencies and better colors. Reds and blues are typically the most difficult to produce, Cranor says. Purple, made from a combination of three dyes, is the most intractable color of all. "Green and yellow," he notes, "are a piece of cake."
There are those who would claim that all this cold scientific knowledge strips away romance. Now that the secrets of chemiluminescence are revealed in stark detail, has the magic disappeared? Hardly. Standing in the dark, green light stick in hand, for me the thrill of the glow remains undimmed. Elizabeth Wilson
Chemical & Engineering News