That pesky second law continues to bedevil abiogenesis

Thermodynamics of the Origin of Life
Brian Miller

Popular articles on origin-of-life research have often portrayed the field as constantly advancing and quickly converging on a purely materialistic explanation for the first autonomous cell. However, the true state of affairs is far grimmer. The highpoint of optimism was actually in the 1950s when Stanley Miller’s experiment produced several of the building blocks of life. Unfortunately, geologists later recognized that the gases used in that experiment did not match the chemistry of the early Earth. When redone with those actually present, the experiments produced  few building blocks in sufficient quantities to contribute to origin-of-life scenarios. Moreover, close examination of modern attempts to synthesize life’s precursors reveals that the  formation of many of them on the early Earth in significant quantities is virtually impossible. However, the more fundamental challenge of thermodynamics dwarfs even these hurdles.

Common arguments about the origin of life have traditionally focused on the unlikelihood of life forming by chance. Perhaps most famously, physicist Fred Hoyle calculated the probability of a cell coalescing to be roughly 1 part in 10 to the power of 40,000. He compared this probability to the chances of a  tornado plowing through a junkyard and assembling a jet airplane. The concept of probability is closely linked to that of entropy, since probability is proportional to the number of configurations (N) in which some state could occur, and  entropy is proportional to the log of N. As an example, the number of ways water molecules can arrange themselves in the solid state is much smaller than the number ways in the liquid or gas states, so ice is the state with the lowest entropy. Due to this connection, the probability argument is often restated that nature tends to move from states of lower entropy to higher entropy, which simply means that nature moves towards states which are highly probable. This tendency is known as the second law of thermodynamics.

Such arguments sound impressive at first, but they are actually incomplete. Nearly all researchers recognize that the first cell could not have come about by chance. They instead believe that some physical processes helped to beat the odds. As an analogy, one could never role one thousand sixes in a row with fair dice. However, if the dice were loaded, that outcome could be quite likely or even close to guaranteed. Analogously, some systems do, in fact, naturally move from states of higher entropy to those of lower entropy (i.e., seemingly low probability) if the lower-entropy states are highly biased to occur. Such a bias is created by a second driving tendency. Namely, nature tends to move from states of higher energy to those of lower energy. For instance, rocks roll downhill, since lower altitude corresponds to lower gravitational energy. Likewise, molecules of water attract each other, so ice is a lower energy state since the water molecules are on average closer to each other than in the liquid or gas states. At low enough temperatures, this attraction overcomes the tendency to move toward higher entropy resulting in water freezing.

However, even in these cases of locally decreasing entropy, the second law of thermodynamics is not violated, for the changes are always exothermic — heat is released. The heat leaving the local system (e.g.,  a cup of freezing water) and entering the surrounding environment increases the latter’s entropy by an amount greater than the entropy decrease of the local system. Therefore, the total entropy of the universe increases. The problem for all theories of origin of life now becomes quite evident. The simplest functional cell compared to its most basic building blocks has both lower entropy and higher energy. And, natural systems never both decrease in entropy and increase in energy at the same time. Such an event would be like rolling countless sixes in a row when the dice are strongly loaded against rolling even one. Therefore, the origin of life through purely natural processes would seem as implausible as water on a hot summer day spontaneously freezing or a river flowing unaided uphill for thousands of miles.

Physicists and chemists often combine entropy and energy (or enthalpy) together into what is called the  free energy of a system. The change of free energy is always negative for spontaneous changes  (e.g., wood burning or ice melting in summer), and it directly relates to the total increase in entropy of the universe. The challenge for the origin of life is then explaining how billions of atoms could spontaneously come together into a state of significantly higher free energy. The chance of such an event happening through thermal fluctuations was calculated by biophysicist Harold Morowitz to be less than 1 part in 10 to the power of a hundred million This number comes directly from an estimation of the free energy of life, and the free energy is a path-independent function. Therefore, this maximum probability estimate is not meaningfully dependent on the theoretical route that led to life (e.g., protein-first versus RNA world) or on the number of steps involved. The probability is always essentially zero.

At face value this thermodynamic analysis for the origin of life would seem to negate any possible materialistic solution to the problem. However, theorists have long recognized one remaining loophole. Morowitz’s calculations assumed that the system was in a state near equilibrium. For instance, some lightning bolt could have struck a pond of prebiotic chemicals at night causing atoms to come together into different configurations. The pond would then quickly settle back into a calm state where the temperature, concentrations, and other variables would remain fairly uniform. However, many argue that the origin of life took place in a system strongly driven away from equilibrium, such as a pond subjected to intense sunlight or the bottom of the ocean near a hydrothermal vent flooding its surroundings with superheated water and high-energy chemicals. Such settings are commonly referred to as non-equilibrium dissipative systems. Their common characteristic is that classical thermodynamics breaks down, so the previous analyses do not completely hold. Instead, principles of non-equilibrium thermodynamics must be applied, which are far more complex and less well understood. Moreover, the energy from these outside sources is hoped to enable the free-energy barrier to be overcome.

However, such appeals to non-equilibrium systems do little to solve the basic thermodynamic problems. First, no system could be maintained far from equilibrium for more than a limited amount of time. The sun is only out during the day, and superheated water at the bottom of the ocean would eventually migrate away from any hydrothermal vents. Any progress made toward forming a cell would be lost as the system reverted toward equilibrium (lower free energy) and thus away from any state approaching life. Second, the input of raw solar, thermal, or other forms of energy actually  increase the entropy of the system, thus moving it in the wrong direction. For instance, the  ultraviolet light from the sun or heat from  hydrothermal vents  would less easily form the complex chemical structures needed for life than break them apart. Finally, in non-equilibrium systems the differences in temperature, concentrations, and other variables act as thermodynamic forces which drive heat transfer, diffusion, and other thermodynamic flows. These flows create  microscopic sources of entropy production, again moving the system away from any reduced-entropy state associated with life. In short, the processes occurring in non-equilibrium systems, as in their near-equilibrium counterparts, generally do the opposite of what is actually needed.