So far we've been talking about thermodynamics the way that physicists and engineers do. And rightfully so. They did all the heavy lifting that gave birth to the field of thermodynamics. But there is a whole other world of thermodynamics: the chemist's thermodynamics.
In the late 1800s, an American scientist by the name of Willard Gibbs realized that the laws of thermodynamics not only applied to heat engines and helium balloons, but also to chemical reactions.
His insight went even further than that, though. He hypothesized that the laws of thermodynamics could predict the future—they could be used to determine whether a chemical reaction would happen or not. As you can imagine, chemists thought (and think) this is Kind of a Big Deal.
The First Law of Thermodynamics remains the same in chemical thermodynamics; the total energy of the universe stays the same, regardless of what happens. Y'know, energy is conserved. The energy might leave whatever system we're looking at, but not the universe. Though that would be undeniably cool.
The gist of the Second Law remains the same, but it's worded a little differently. It says that all spontaneous processes increase the total entropy, or disorder, in the universe.
Physicists said that heat will always flow from a warm object to a cold object, not the other way around. That's because giving a cold object some heat increases its disorder, or entropy. Kind of like giving a kid a bunch of Legos also increases the amount of disorder in the area. A warmer object's molecules move about more quickly than a colder object's molecules, so the warm object's molecules are in greater disorder.
Because of the chemical reactions they observed, Gibbs and his students added two new ideas to our understanding of entropy: enthalpy and Gibbs' free energy. Don't worry, they're not as '70s as they sound.
Chemical reactions involve the breaking and forming of bonds between atoms. These reactions either take up heat or give off heat. Enthalpy, H, is a measure of the change in heat when a chemical reaction occurs. It tells whether the reaction is endothermic (and so requires energy or heat input), or whether it is exothermic (and gives off energy in the form of heat).
Whether chemical reactions are spontaneous (like a pile of fireworks sitting in the middle of a forest fire) or not spontaneous (like those same fireworks sitting in your freezer) depends on both the change in entropy as the reaction occurs, and the change in enthalpy. In other words, it depends both on whether disorder is increased or not and on whether the reaction requires energy or releases energy. Gibbs' free energy, ΔG, combines the change in entropy and the enthalpy of a reaction, and tells whether or not this reaction will occur spontaneously. It's a chemical soothsayer.
If ΔG is negative, the reaction will occur spontaneously without any energy added. If ΔG is positive, the reaction is not spontaneous because it requires additional energy, and if ΔG is zero, the reaction is in equilibrium, neither giving off nor taking in energy. It's zen.
We've talked a lot about heat and temperature, but we still haven't really explained why 90 °F feels hot and 50 °F does not. Sure, at 90 °F objects have more internal energy than at 50 °F. Molecules are zipping about a lot faster, but what does the internal energy of the air have to do with how we feel? And how do zippy air molecules make us sweat buckets?
The answers to all of these questions lie in the Second Law of Thermodynamics. Heat always flows from warm objects to cold. If a warm object is next to a cold one, the heat will shimmy about between them until both objects are at the same temperature. The universe, essentially, strives for everything in it to be at the same temperature. Obviously, it has a long way to go to achieve that goal.
Our human body temperatures average a comfortable 98.6 °F. If it gets too much higher than that, processes within the body start going terribly, terribly wrong. Proteins denature and cells die. It's kind of like being cooked from the inside out. So, not our idea of a good time.
Being feverish is a problem even on a molecular level. See, all of the normal chemical reactions that keep us alive raise our body temperature, so our bodies works very hard at kicking out all of that extra energy—even when we're not home sick in bed.
On a 50 °F day, there is a large heat gradient between our internal body temperature and that of the outside. All that heat generated by our body can quickly escape down that steep heat gradient (as if the heat were "rolling down" a hill) through our skin by radiation and conduction. On a 50 °F day our bodies are like a hot reservoir surrounded by much cooler air.
On a sticky summer day, when the air temperature approaches the value of our body temperatures, it's harder for heat to escape our bodies. On a 100 °F day, heat is actually flowing into the body. Now we can't dump our extra heat through radiation and conduction. The only way to keep cool is to sweat. That sweat can then evaporate from the skin, taking with it some unwanted heat. In the Southern U.S. or in the tropics, even sweating isn't enough (because of the high humidity reducing evaporation), and that's why fans were invented.
How does evaporating sweat cool us down? Is it any surprise that the answer is "thermodynamics"? Think about it: what does it take to turn liquid sweat into a gas? It's a phase change, and those things need energy. We have to get those sweat molecules moving faster and faster, and the energy for this comes in the form of body heat. Body heat is used up in evaporating the sweat. With every bit of sweat we evaporate, our bodies are able to lose some unwanted heat.
That's why 90 °F feels hot and uncomfortable: because our bodies have to work harder (read: sweat) to maintain normal temperature. On a 50 °F day, dumping heat is a cinch, with no need for sweating, fans, or distress.
After spending all of this time discussing the laws of thermodynamics, we're going to throw a curveball. There's pretty good evidence that these thermodynamic laws only apply to big stuff like heat engines, power plants, and chihuahuas. Wait, how big are we talking here?
When objects get very small, like molecular sized, the laws of thermodynamics aren't as accurate3. For a long time scientists didn't know that, because they couldn't study the behavior of single molecules. Turns out, getting a bigger magnifying glass just doesn't cut it.
For example, for a long time scientists couldn't study exactly how the polymerase that replicates our DNA works. They could study the average behavior of the millions of polymerases in our personal test tubes (bodies), though, and that average behavior obeys the laws of thermodynamics.
Because of the relentless march of science, we can now study a single DNA polymerase or a kinesin (a little motor that acts like a tow truck, moving things around our cells). Check out this video that shows how a virus can stuff DNA into itself. The giant leaps in technology that now allow us to study single molecules allow us to build detailed models of what goes on inside of our cells.
Why do we care about all of this? As long as the hundreds of polymerases in our test tube behave like the thermodynamic laws predict, who cares if the individual polymerases are a little more fickle?
We care because the things we are able to observe are getting smaller and smaller. People want to build molecular motors that can be turned on and off by light particles, and quantum computers that would be able to process information much faster than the digital computers we work with today. We need to understand the thermodynamics of tiny things to develop these of nano-technologies. This is where quantum dynamics comes into play. And plays with our minds, too.
Professor Jonathan Oppenheim at University College London recently found that for very tiny objects, processes are nearly always irreversible. There are no efficient Carnot engines on the nano-scale. His research suggests that any heat engine or micro-powerplant that technologists want to build will be far less efficient than its bigger counterparts. Talk about a major bummer. This just goes to show we are never done learning and discovering. And sometimes getting bummed out along the way.