Prof. Harry Gray the six-foot three California Institute of Technology inorganic chemist, National Academy of Sciences member, and likely candidate for a Nobel Prize sometimes lectures as a horse, or as a leopard. His students retaliate by gluing his hand to his telephone, turning the lecture hall seats backward, and on one memorable occasion binding a nude centerfold of him into the Caltech catalogue.

"It was tastefully done," chortles Gray. "To this day only a select few know if it was my body or not."

Students have turned Gray's office into Lavoisier's cabinet, complete with cobwebs, and into a high-tech miniature golf course with blow drier tees and water hazards. (Though this last venture put Gray out in the hall for a month, he graciously concedes that the golf was "very elaborate, very exciting.")

Gray is himself a prankster, and his students and postdoctoral fellows follow suit. Yet many go on to become luminaries in the world of science. Seventy hold top academic job. As for his department, a recent Science News "citation impact" report bore the unequivocal headline: Caltech Chemistry No.1 in World.

Trained as a classical inorganic chemist, gray today studies the tiny inorganic bits ("where the action is") in living molecules. This is "bioinorganic" chemistry, a new field Gray helped invent.

"Organic molecules contain carbon, hydrogen, and oxygen," he explains. "But we now know that practically everything organic also contains inorganic bits. Each mammoth hemoglobin molecule, for instance, has within it four tiny iron complexes that bind to oxygen molecules and carry them around."

Biochemists have long recognized the presence of trace metals in organic substances, but until Gray became interested in the color of blood, few realized just how critical the metal ions' structures are to the processes of life.

What makes a kid do science?

Harry Gray was born 57 years ago in Woodburn, Kentucky, close to the Tennessee border. In his day, Woodburn has some 200 residents; today, it has nearer 50. (The town recently voted 23 to 0 to discontinue a tax supporting its one streetlight.)

Gray's grandfather was a blacksmith, his father a high school teacher, his relatives' farmers and all of them natural scientists. "They constantly experimented," he says of his uncles, one of whom later left farming to become a physicist. "They built ham radios, they fixed machinery, and had an enormous ability to do practical science."

From the time he was 10, Harry, too, began experimenting, "driven toward chemistry by some mystical something.

"I used to visit my mother's mother on the farm," he explains. "More than anything else in the world, I loved my grandmother. But when I was 10 she got cancer and went to the hospital in Nashville. After suffering a great deal, she died."

Appalled by the implacable finality of cancer, young Gray determined to do something about it.

"At 10, chemistry seemed to me to be a way to express my indignation not that I thought chemistry could cure anything, but I felt that it was something I could do. I began writing notes to my mother saying 'I'm going to be a chemist. I'm going to learn how to do chemistry and do something good with it.'

"Of course, I had no idea what I was talking about. I started out in inorganic chemistry, and the work I've done since has only a distant connection with biomedicine. But somehow, when my grandmother died, I knew I had to pursue science."

Education and color

"From the beginning I was interested in mixing chemicals to make explosions and flames, but chiefly I was interested in colors what made them, and why some compounds were one color and some another. It's the one theme that's stuck with me throughout my life 20 years out of high school; I was trying to explain why proteins with copper in them are a rich, deep blue.

"By junior high I had developed an elaborate chemistry laboratory. A company in Chicago sold me chemicals, and in my basement, I mixed them and looked at colors. By high school I had a lot of practical knowledge, but no explanations. My teachers were wonderful, dedicated, brilliant women (whom, frankly, the world was taking advantage of). The curriculum had very little science but was otherwise classical and rigorous, and these women really shaped me."

In addition to running the school's chemistry lab and being an excellent student, Gray was the class "techie" who did special effects at school functions.

"I got straight A's, but the girls didn't want to have much to do with me. Then on year, determined to attract a certain girl, I began getting zeroes in English class. My teacher, Mrs. Craig, saw immediately what was going on. 'Harry,' she said, 'I know way you're flunking this course, but that girl won't pay attention to you no matter what you do. Some day you will find girls who like you because you do know something. Right now, you should have some self-respect. If you've got any guts, you'll do perfect work from now on, and I'll give you a B.' She really impressed me, and I did."

On his senior day, Gray set off a smoke bomb that emptied the school for two days. "I scaled up," he blushes, abashed still at the memory, "and it got a little out of hand."

Up from Kentucky

At Western Kentucky University, three dedicated chemistry teachers took Gray in hand and gave him the solid theoretical background he needed to match the practical knowledge he had acquired in his basement.

"I had no idea why the silver ion reacts with ammonia," he recalls, "though I had mixed them together often enough. When you add copper ions to the ammonia, you get 4 ammonias bonded to copper. But why is it 4 with copper, and 6 with cobalt? I knew some formulas, and I knew the colors they formed, but when I entered college, I still had no idea why."

In a physics class Gray met his mathematician wife-to-be Shirley, and when he graduated at 21, they were married. It was time to leave Kentucky. "Chemistry," he observes, "was my ticket out." Turning down a scholarship to the Massachusetts Institute of Technology, Gray chose instead to do his graduate work at Northwestern University in Evanston Illinois.

At Northwestern, Gray met Fred Basolo and Ralph Pearson, great pioneers of modern inorganic chemistry, who introduced him to reaction mechanisms and crystal field theory and changed his life.

Throughout the 1930s, 1940s, and 1950s, Linus Pauling a great theoretical chemist and probably the world's greatest lecturer successfully promoted his "valence bond theory" that atoms stick together by using hybrid orbitals to make covalent bonds with shared electrons. At the same time Robert Mulliken an equally great theoretical chemist and probably the world's worst lecturer was developing his more complicated "molecular orbital theory," which, in addition to allowing for electron sharing, predicted accurately how light is absorbed in different places and explained how light-struck electrons are promoted to excited states.

By 1960 inorganic chemists interested in the color of compounds were realizing that Mulliken's molecular orbital theory largely ignored by the inorganic crowd for 25 or 30 years was superior to Hans Bethe's crystal field (or ionic model) of chemical bonding, according to which electrons are completely localized in charged ions.

"Mulliken's was a realistic bonding model, not a simplistic one where all the electrons are stuck on the atoms," says Gray. "With it, you could go toward Pauling's idea of sharing but still have a model that explained color."

Copenhagen, crystals, and the color of everything

In 1960 the Grays arrived in Copenhagen on a National Science Foundation postdoctoral fellowship. During the year they spent there, Shirley gave birth to one of their three children while Gray learned pretty fair Danish and made major advances in molecular orbital theory.

"I grew up in Copenhagen," he remembers with pleasure, "and my mentor was Carl Ballhausen. He taught me that my own expectations were not nearly as important as designing the experiment, doing it, getting results, and then moving on even if it meant discarding my original idea entirely. I learned that being wrong could be an absolute gain, even a pleasure, because then I finally knew something and could go on."

When Gray arrived in Copenhagen, both physicists and inorganic chemists studying crystal field and molecular orbital theory were dealing with dull compounds simple crystals and metal ions in water. But by year's end, Ballhausen's knowledge of theory and Gray's strong experimental background had produced phenomenal results. Together they demonstrated how Mulliken's molecular orbital theory called "ligand field theory" when applied to metal based compounds could be applied to complex inorganic molecules important in the real world.

"With molecular orbital theory, I could look at a ruby and tell from the structure of the crystal why it was red. A ruby is chromium aluminum oxide, with the oxygens surrounding each chromium in an octahedral structure. This structure, or ligand field, determines how light is absorbed.

"Ligand field theory was marvelously simple, yet it enabled us to calculate exactly where metal ions in a substance will absorb light. For ruby, it predicted light absorption peaks at certain wavelengths and there they were. It was incredible! You cannot believe how exciting it was that, all of a sudden, electrons jumping up and down in orbitals in a ligand field could explain the absorption spectra of complex inorganic molecules.

"Most annoyingly, Pauling paid no attention to us, and his third edition of The Nature of the Chemical Bond devoted only a page to ligand field theory. But our discovery had finally brought Mulliken's theory to the fore, and he was delighted.

"Ligand field theory was revolutionary, and it became the rage of the sixties," says Gray. "Inorganic chemists began falling all over themselves explaining the magnetism and colors of everything."

New York, New York

In 1961 Gray got his first real job. As assistant professor of inorganic chemistry at Columbia, he began teaching ligand field theory to freshmen,

"Pauling's model was in the textbooks," he remembers. "But I said, "Look, with ligand field theory the kids are going to understand all about color and energy levels and get turned on.' And I was right. I had some brilliant freshmen, and they loved it."

More of a New Yorker than New Yorkers, Gray loved it, too. In his second year there, he began giving inorganic chemistry lectures to biologists and biological chemists at Rockefeller University. It was an eye-opener.

"Chris Walsh," he remembers, "a Rockefeller student who is today head of a department at the Harvard Medical School, would come up to me after class and say, "Look, if you can explain the colors of simple iron compounds, you can probably explain hemoglobin and myoglobin [iron-based proteins that transport and store oxygen]. Harry, apply what you're doing to biological molecules. You can make a difference.

In 1963, Gray was named associate professor at Columbia and, in 1965, full professor. But by then he had decided to study proteins (for which he needed biologists) and the reactions of molecules energized by light (for which he needed photo chemists). Caltech offered the best of both, and in 1966, he left New York for California.

Caltech and new bioinorganic chemistry

With the structure of hemoglobin already known, Gray decided to work on hemerythrin a reddish purple protein founded in the blood of worms that, like hemoglobin, is iron-based.

"The iron in hemerythrin is more like inorganic iron, without the large home group of organic stuff surrounding it," he explains. "This was closer to something I thought I could understand."

But having traded in nice, neat crystals for worms, Gray's first problem was obtaining biological samples. "This was horrifying stuff," he recalls, "bleeding worms, cutting them up, putting them on funnels with glass wool, and letting the blood drip through. But I had to do it to isolate proteins."

Caltech physicist Jim Mercereau provided sensitive machines (superconducting quantum mechanical interference devices, or SQUIDs), so that Gray and his coworkers could measure the magnetic properties of metals in proteins. "If two iron atoms are close together," explains Gray, "their electron spins tend to cancel so that the thing is less magnetic than it would be if the irons are isolated. We found that two irons are, indeed, close together in hemerythrin."

After worm blood, Gray began puzzling over the color of copper. New copper is gold, copper roofs turn green over time, copper sulfate is a pale blue, and, the real mystery, certain copper proteins are a deep, rich, azure blue. Inorganic chemists had struggled long and hard to make the deep blue copper compound, but as of 1969, when Gary entered the fray, none had succeeded. Spectroscopy, he decided, was the key to this Holmesian search to find the geometrical arrangement of atoms surrounding copper in azure blue, bioinorganic protein molecules.

"We figured there had to be a sulfur and a couple of nitrogens in a certain configuration around the copper," he says, lightly summing up years of work. "It turns out you can't get this configuration in a simple compound, but we pretty much figured out how you could have this structure in a protein" (findings X-ray diffraction has now largely confirmed).

During the next several years, Gray and his Caltech colleagues used color and magnetism to determine the structure of many bioinorganic molecules.

Life and death by electron transfer

Then Gray began wondering how bioinorganic molecules work.

He knew that proteins with iron or copper transfer electrons. Worm blood just ferries oxygen around. But the beautiful, blue, copper proteins not only transfer electrons to other proteins, some also reduce oxygen to water. Bioinorganic electron transport systems, furthermore, are virtually perfect, with proteins delivering electrons to the right place every time.

"If they didn't," says Gray, "we'd all be dead.

"First, we eat lunch, which is organic stuff full of electrons. Then chemical reactions very carefully extract electrons from the food and move them along a molecular chain to their destination. Some molecules barely hang on to their electrons, while others hold on tight. But in the end, lunch's electrons react with oxygen to give water and carbon fragments, which are then broken down into carbon dioxide."

The energy liberated by this is stored in the form of adenosine troposphere bond breaks apart to make stronger bonds and emits a small squib of energy.

"Your body can't afford explosions," explains Gray. "Food is the body's equivalent of gasoline, but it doesn't combine with oxygen like fuel in your automobile engine. With ATP, every thing is nicely tuned no BANG! BANG! Reactions. If the food in your body reacted directly with oxygen, you'd burn up!

"But if electron transfer in cells cannot go too fast, neither can it stop. Look at cyanide. It is a terrific ligand, and once it binds to the iron in cytochrome oxides [the protein that reduces oxygen in the body], oxygen cannot bind there and receive electrons. A few seconds of that, and you're dead."

Jumping electrons

In biological systems, electrons jump from metal to metal all the way down the electron transport chain. And as Gray discovered, they really do jump.

For many years, inorganic chemists thought that, to exchange electrons, ions had to be close enough to overlap. But if that were true (as it is in inorganic chemistry), how was the process in biological systems so perfectly controlled? And how did the few metal atoms in proteins, surrounded by thousands of atoms, come close enough to each other to overlap?

Yet inorganic chemists Gray included firmly believed that electrons could jump no more than 2-4 angstroms, that is, over a single atom. "There was a lot of classy work," he remembers, "showing that, if you wanted to get electrons from a ferrous to a ferric iron fast, you darn well had to get them close together."

Biologists, on the other hand, believed electrons could jump huge distances fast, some 10-20 angstroms, over 10 atoms of more.

Everyone agreed that electrons ended up long distances away from their starting points, but how? By contorting the intervening material or jumping over it?

The 'fixed distance' experiments

"We chose cytochrome-c, a common iron protein found in the respiratory systems of mammals, because its structure is rigid and not easily deformed, and because anything fixed to its surface in a certain region has more than 10 atoms between it and the center and cannot, therefore, be closer than 12 angstroms to the iron atom located there.

"We then tried to attach a metal to the outside of the cytochrome-c molecule," he continues, "by bonding the metal to an amino acid. We tested countless different metals. We used atomic absorption analysis to check whether the metal was present and peptide mapping [cutting up the protein with trypsin enzymes and taking out little pieces] to isolate the bit with the metal on it.

"For eight years we had no success, but we refused to give up."

Finally, Gray's group tried ruthenium, and in 1981 Kathy Yocom, one of Gray's students, succeeded in isolating, characterizing, and proving the exact spot where the ruthenium was bound. Eight years has given them a suitable molecule and a technique for injecting an electron and watching it jump.

The answer was immediate and absolute: Electrons jump.

"I was ecstatic. Our whole group was ecstatic," says a gleeful Gray. "The best thing that ever happened to me was to be wrong about this, because it is so much more exciting that way. We had expected this particular jump to take hours, and it took about a few thousands of a second perfectly compatible with the times we knew for many electron transfers in biological systems."

There is only one way for electrons to travel that fast and accurately across long distances that is, by jumping from one atom to another down a chain until they reach their destination. Gray and his students fixed probes at various points on the surface of the test molecule to show that electrons on different routes traveled at different rates. They also showed that these rates were predictable.

A tunnel with a break that prevented electrons from getting through was dubbed "cold," while quick and easy routes were "hot." Control of the whole process, according to Gray, depended on the chemical structure of the atoms along the tunneling pathways.

"There are bonds that you jump through easily," he explains, "and others you apparently can't jump through at all. There are places where the rule that an electron travels across "five atoms in a trillionth of a second' which is the general rule we go by simply does not work. If there is a cold spot in those five atoms, or in any one of them, they can slow down the flow or even stop it dead."

End of the rainbow

"I spent 10 years working out structures from color," summarizes Gray," and another 20 working on the electron jump problem. We are now trying to understand how electron tunneling where electrons travel down a chain of atoms to their final destination works. Theoreticians are calculating specific 'tunneling paths' and producing electron transfer maps, while experimentalists like myself are trying to check these maps."

Gray expects this stage to go on for another few years. Then, once scientists understand the process, they will try to control it.

"Controlling electron flow in bioinorganic systems has vast implications for mankind," he goes on. "We could soon be making really superior solar energy storage devices to fuel cars and heating systems. Molecular, perhaps even bioinorganic, devices could replace solid-state silicon chips for storing and manipulating information in supercomputers.

"There's a big payoff coming in this field," predicts Harry Gray, radiating equal parts of enthusiasm and confidence. "And it's not far off."

Editors Note: Caltech's R.A. Marcus, who has done much of the theoretical groundwork in the field of electron transfer and calculated many of the systems measured experimentally by Harry Gray's team, received the Nobel Prize for chemistry in 1992.