American Scientist

One afternoon in 1842, in the town of Walsall in the heart of England's industrial midlands, two young men stood by a canal, watching a lock fill with water. The rising water lifted a barge crammed with valuable trade goods, one small step up on its climb to some unknown industrial destination. The two men mused upon this ingenious use of power, this impressive demonstration of the simple technology underpinning Victorian Britain's industrial dominance.

The two men were brothers. One was James Thomson, a shipbuilder's apprentice later to become Professor of Engineering at Glasgow University. The other was James's brother William, destined for an even grander career. William's sojourn as Professor of Natural Philosophy—also at Glasgow—would span half a century and include fundamental contributions to an astonishing range of sciences and technologies, from the transport of fluids to the design of ultrasensitive telecommunications. William Thomson would ultimately be ennobled by Queen Victoria, becoming Lord Kelvin of Largs.

December 2007 sees the centenary of Kelvin's death. That early curiosity about energy, shared with brother James as they stood by the Walsall canal, was just the beginning of Kelvin's part in the most significant transformation of physical science since Newton. In tandem with others, such as French engineer Sadi Carnot, German physicist Rudolf Clausius, and English experimenter James Joule, Kelvin developed the science of thermodynamics: the fundamental understanding of the nature of heat, energy and temperature.

The key to the Industrial Revolution and to Britain's dominance in the 19th century was the use of energy. But what fascinated the Thomson brothers that day by the Walsall canal was the potential waste of energy. What became of the power lost when water cascaded over the gates of the lock instead of helping lift the barge? What rules determined how much power was wasted—in this or any other of the myriad technologies on which industry relied?

This question, it turned out, was not just the key to industrial efficiency, but to understanding the nature of all energy transformations—of everything that happens in the Universe, essentially. In pursuit of an answer to this question, Kelvin in Glasgow and Rudolf Clausius in Germany developed perhaps the most fundamental physical law, the Second Law of thermodynamics, which states that every exchange of energy is wasteful; perfect efficiency is impossible.

Whereas Kelvin developed many of the concepts behind thermodynamics, it was Rudolf Clausius who expressed the Second Law in mathematical terms. Clausius introduced a quantity he named entropy, which describes the state of order or structure in a system. According to the Second Law, entropy or disorder (a simpler if not entirely accurate synonym) increases with every energy transaction. Hence, entropy is a sort of dark twin to energy, one that contrasts energy's constancy (the First Law states that energy is conserved) with its own inevitable rise.

Kelvin and Clausius thus presided over the first thermodynamical revolution. The force behind it was a practical one: the need to harness energy for industry. Its dogma was the triumph of entropy and disorder. Yet a hundred years after Kelvin's death, research in thermodynamics is questioning this message of chaotic doom. Some scientists have begun to step beyond the major limitations of 19th-century thermodynamics to explain why our Universe seems characterized not by spreading disorder, but rather by a fantastic degree of structure, complexity and creativity.

The Shock of the New

Kelvin and Clausius had grasped the essence of energy, heat and temperature—with, however, certain caveats. Kelvin's thermodynamics was a science of equilibrium: fine for systems going from one stable state to another, but saying little about systems that lacked this stability. It was also an isolated science: Kelvin's theories applied only to so-called closed systems, which are, by definition, unaffected by their surroundings (such as a steam engine that used heat to do work). And it was a "science of the large," one that described practical hunks of material such as boilers. Early thermodynamics was adequate for predicting the behavior of billions upon billions of atoms, but no good for tiny systems—such as the complex molecules found in living things.

Modern thermodynamics is all about stepping beyond these limitations to understand processes and systems that are far from equilibrium, inextricably tied to their surroundings, and at least a million times smaller than industrial-scale engines. Consider one such molecular engine: the protein. In converting stored chemical energy to useful work—such as transporting cargo, catalyzing chemical reactions or pumping ions across membranes—proteins do for life just what the big engines of Kelvin's day did for industry. As current research in thermodynamics conquers the limitations of 19th-century theory, it is faced with one of the great challenges of modern science. The new goal is a second transformation more profound, perhaps, than Kelvin's: to go beyond the thermodynamics of the inert world toward the thermodynamics of life.

Thermodynamics Under the Microscope

Proteins are incontestably beyond the limits of Kelvin's thermodynamics, for several reasons. Proteins are not at all "equilibrium" machines; the inside of a cell is a profoundly heterogeneous place of concentration gradients and complex spatial structures. The size of a protein also limits the reach of Kelvin's laws: A typical protein molecule contains a few thousand atoms and is only tens of nanometers long. Precisely because of this scale, proteins are inextricably part of the chaotic molecular sea that surrounds them—one that could never be called a closed system.

Fluctuation is the key to understanding thermodynamics in small systems. The Second Law as developed by Kelvin tells us that, on average, entropy must always increase. But when a system such as a protein engine is made from a small number of components—and compared to the billion billion billion molecules of gas in a steam engine, even a ten-thousand-atom protein qualifies as small—you can't ignore fluctuations about that average. The water molecules surrounding the protein are subject to Brownian motion, which sends them crashing into the working parts of the protein engine. Not surprisingly, the energy supplied by the engine tends to fluctuate by an appreciable amount.

Over the past decade and a half, scientists have developed theories about thermodynamics at a microscopic scale. Christopher Jarzynski at the University of Maryland, Denis Evans at the Australian National University and Carlos Bustamante at the University of California, Berkeley, have all studied the application of the Second Law to the world of microscopic engines. Through experiments that tracked the motion of individual molecules, Bustamante's team in 2002 watched how a stretched RNA molecule relaxed back to its equilibrium shape. Two years later, Evans' group used a laser trap to tug on tiny particles in the grip of Brownian motion. Such experiments provide spectacular views of microscopic thermodynamics in the real world.

The primary conclusion of such studies is that biological engines such as proteins—not to mention manufactured engines on the same nanoscale—are indeed caught in a world of inevitable fluctuation. Understanding their biophysical function means figuring out how such engines not only cope with fluctuations, but use them to maximize efficiency. Scientists such as Alan Cooper of Kelvin's own University of Glasgow and Hans Frauenfelder at Los Alamos National Laboratory have spent years investigating this topic. Data from Frauenfelder's group show that a protein's fluctuating shapes form a complex hierarchy that may explain the versatility and efficiency of such machines in a difficult environment.

Understanding these microscopic protein engines at a thermodynamic level would mark a huge achievement in science and, potentially, a revolution in medicine. Imagine curing disease by reengineering the microscopic workings of the cell rather than by prescribing some chemical cocktail. Of course, such ability raises serious ethical dilemmas. However, ignorance is no protection: Our growing understanding of microscopic thermodynamics is the key to safe, ethical use of nanotechnology.

In Kelvin's time, thermodynamics enabled the industrial revolution—a historical period that produced its own ethical and social questions. Similarly, microscopic thermodynamics will enable a new revolution and prompt more discussion about the appropriate uses of technology.

Escaping Entropic Doom

Kelvin expressed the Second Law at its cruelest with his prediction that the universe was headed for a grim end called universal heat death. Because perfect efficiency is impossible, wasted energy—another way of describing entropy—always increases. The dissipation of useful energy as heat, taken to its logical conclusion, means that all useful energy in the Universe will eventually be frittered away.

But the Second Law is a law of equilibrium—a law of closed systems. What if the important processes in our Universe were overwhelmingly not at equilibrium, were not the products of isolated, closed systems?

Indeed, this is precisely the case: Equilibrium is extremely rare in the natural world. Take the earth's atmosphere—a system profoundly out of equilibrium. The air we breathe is a potentially unstable mix of oxygen, nitrogen, water vapor and so on, and could quickly disappear in a flash of furious chemistry. Yet it has not done so. This is no chance fluctuation: the presence of life and the continual influx of solar energy maintain this out-of-equilibrium state. By means of non-equilibrium thermodynamics, Earthly matter has formed a complex, yet highly ordered and stable system.

Kelvin's thermodynamics, limited to closed systems at equilibrium, was thus the tip of the iceberg when it comes to energy in the Universe. But scientists and philosophers have spent much time contemplating the rest of the berg in the form of non-equilibrium, open-system thermodynamics. Belgian chemist and Nobel laureate Ilya Prigogine suggested that systems away from equilibrium, and especially those fed with a steady supply of energy, are not bound to the disordered fate implied by the Second Law.

In open systems, entropy is not so much a harbinger of doom as a currency for creativity, because in an open system, local entropy can fall as well as rise. True, any decrease must be paid for by a greater increase somewhere else, but Prigogine and his followers argue that the swap of order here for disorder there reveals the inherent creativity of matter that is far from equilibrium. A continual supply of fresh energy enables order to arise from disorder; the phenomenon of so-called self-organization, which cuts across all scales, ultimately results in a creative universe.

Open Universe, Open Life

The thermodynamics of open systems remains a controversial area, sparking debate reminiscent of similar controversies over complexity and the emergence of life. The similarities are not accidental: All these ideas pit the spontaneous creation of complexity or order against the Second Law's mandate for disorder.

But debate alone will not solve the problem. Perhaps the best way to learn more about non-equilibrium and open-system thermodynamics is not to formulate grand theories, but to study their consequences in the form of palpable objects, actual pieces of matter. Kelvin's profound ideas about energy led to a fundamental revolution in modern science but were firmly anchored to a solid reality in the industrial engineering of Victorian Britain.

Which brings us back to microscopic engines: the most interesting objects in which the two themes of modern thermodynamics—microscopic scales and open systems—join. Although studies of individual proteins are important foundation stones, the cell depends on millions of molecules in a complex network of machines, their functions interlocked across a range of scales. Such interplay is possible precisely because these living engines are open to fluctuations and not isolated from their environment. It may be that the complex functions of matter that we call life are nothing more than this multiscale interplay of engines, a network through which energy is transformed again and again, as microscopic machines swap and shift matter—manipulate entropy—in a thermodynamical cycle the likes of which Kelvin could hardly have imagined.

If so, then the theory that accommodates living engines will require both a thermodynamics of microscopic matter, and a thermodynamics of open systems. Current research is providing progress in each of these, but the challenge may be to match them together, to join the palpable chemical reality of protein engines to the heady concepts of non-equilibrium thermodynamics and self-organization. That may lead us, finally, to the second revolution—the thermodynamics of life.

Kelvin's theory of thermodynamics was only one product of his long, intensely curious life. At retirement, he declared that despite half a century of work he felt he understood little more about the nature of the physical world than he had all those years ago, staring down at the water of the Walsall canal. Now, a century after his death, the science of thermodynamics that Kelvin pioneered is indeed more puzzling, more profound, more tantalizing, more practically relevant, and just plain more fascinating, than ever.