The Cool Power of Evaporation
Nature can do amazing things — if you don’t mind working up a sweat.
In January 1839, the first shipment of a valuable commodity arrived in Sydney. It had been harvested on the other side of the world months earlier, and then shipped in insulated boxes over thousands of kilometres of ocean from Boston. Its arrival brought such relief to long-suffering Sydneysiders that the Sydney Herald compared it to a gift from the gods of Olympus.1 Just as the Promethean gift of fire had long ago provided humanity with protection from the cold, finally there was a counterpart to provide respite from the heat. Ice had arrived in Sydney. Indeed, by 1861, thousands of people worked across the United States in extracting and transporting hundreds of thousands of tonnes of ice all around the world, at a cost of hundreds of millions of dollars in today’s money.
That, until the invention of artificial refrigeration, this peculiar trade was the most viable way of achieving large-scale cooling reflects a profound fact of nature. It is much easier to stay warm than to stay cool. This is because any source of energy can be transformed into heat. Indeed, our ancestors discovered hundreds of thousands of years ago how burning fuels like wood releases immense quantities of heat. Our bodies naturally do the same with the energy we consume from food. This made it possible for even pre-modern humans to live at temperatures that were routinely as much as eighty degrees Celsius below our body temperature of thirty-seven degrees.
By contrast, humans are susceptible to heat exhaustion at temperatures even just a few degrees above body temperature. This is because, unlike heat, coldness is not an entity that can be generated; it is simply the absence of heat. The only way to cool ourselves is to quickly lose heat to our surroundings. If the air around us is hotter than us, this does not happen and so we overheat. To avoid this, we need a heat sink — something that will absorb heat even on a hot day. Ice is one such heat sink, simply because it is very cold and stays cold for a long time. However, as we have seen, it took the might of nineteenth century industry and capitalism to make ice available in warm climes. For most of human and natural history, this was not an option.
Fortunately, there is a natural alternative — sweat. The ability to cool ourselves by sweating is essential to our ability to maintain a constant body temperature in hot weather or while exercising. Indeed, it is so vital that it is thought that humans lost the typical mammalian covering of fur primarily to allow for more efficient sweating. Yet, the cooling effect of sweating is ostensibly quite mysterious. Sweat is not itself cold like ice; it is just as warm as the rest of our body. How then, can it act as a heat sink, leaving our skin feeling cooler? To answer this question, we must explore the science of evaporation.
The Science of Evaporation
While ice only forms under very cold conditions, water naturally exists around us in both of its other two states. Liquid water can of course be found in oceans, lakes, rivers and puddles, as well as inside organic matter like our bodies. Gaseous water — water vapour — makes up a fraction of the air. These two states coexist because they are each favoured by one of two competing physical effects.
On the one hand, water molecules are drawn towards one another by a type of electrical attraction called hydrogen bonding. This leads them to clump together into droplets that are too heavy to remain suspended in the air and so fall to the ground, where they collect together with other fallen droplets. In this way, water vapour can collect into a liquid in a process called condensation.
On the other hand, the effect of entropy pulls water towards the gaseous state. As explained in Order From Chaos, this is because molecules tend to spread out to fill a room because there are overwhelmingly more ways for the particles to be distributed evenly across the room than to be gathered in one place. When water molecules distribute uniformly across space like this, they form water vapour. So, water molecules that are initially gathered into a liquid can also be pulled away and spread across the room to become gaseous in a process called evaporation.
Whenever two objects are pulled together by a force, they lose potential energy (as discussed in How Electricity Powers The Modern World). This energy is transformed to heat energy which dissipates into the surroundings, heating them up. An example of this is a heavy object falling as it is pulled to the Earth by gravity; when it hits the ground it releases energy, as can be observed by the sound it makes and the destruction it can cause. Water condensing is another example; the water molecules release heat as they are pulled together. This means that water releases heat as it condenses, warming its surroundings. Conversely, water molecules pulling apart must gain energy, just as you must use energy to lift a heavy object off the ground. The molecules absorb this energy from the heat of their surroundings. This means that water absorbs heat as it evaporates, cooling its surroundings!
This is why evaporation of sweat cools your skin. The evaporating water absorbs heat from your skin and the air around it. This reduces the temperature of your skin, leaving it cooler. In this way, liquid water — like sweat — can act as a heat sink even without being any colder than its surroundings.
Forever Out Of Balance
Yet, this still leaves an important question — why does liquid water evaporate more than water vapour condenses? We are used to the idea that wet surfaces tend to dry as the water evaporates off them. Yet, as we have seen, the mutual attraction of water molecules is also constantly acting to condense water vapour into liquid. So why doesn’t this condensation cancel out the effects of evaporation, replacing evaporated liquid and releasing heat that counteracts evaporation’s cooling effects?
In reality, condensation and evaporation are always occurring. How quickly each process occurs depends on the relative humidity, or simply humidity. This measures how the amount of water vapour currently in the air compares to the maximum amount that the air can hold. Humidity is relative because warmer air can hold more vapour than cooler air, meaning that the same absolute quantity of water vapour corresponds to a lower humidity at higher temperatures than lower temperatures. However, it works well as a measure of the tendency for water to evaporate or condense.
Specifically, when the humidity is below 100%, there is less water vapour in the air than it can hold, so water tends to evaporate faster than it condenses to increase the quantity of vapour; the reverse happens when the humidity is above 100%. In this way, humidity naturally tends towards 100%, which is called equilibrium. In equilibrium, evaporation and condensation occur at the same rate. At this point, the heating effect of condensation does indeed cancel out the cooling effect of evaporation.
However, the humidity close to the ground is nearly always less than 100%.2 For example, in Sydney, the average afternoon humidity is just 56%. Under these conditions, water evaporates quite quickly and so we can be cooled effectively by sweating. On a particularly humid day, when the humidity is much closer to 100%, evaporation happens more slowly. This is why warm, humid days can feel much hotter than hot, dry days. But even under these conditions, some evaporation still occurs and so cooling can still occur, albeit much more slowly.
However, this still leaves the question — why is the humidity typically less than 100%? Why has the ongoing process of evaporation not long ago resulted in equilibrium? The answer is that water vapour does not just spread out horizontally around a room; it also tends to spread vertically. This means that some water vapour naturally rises upwards higher into the sky. Further above the ground, the air is cooler, meaning it cannot hold as much water vapour. As a result, the rising vapour eventually reaches a height at which it naturally condenses back to liquid droplets, forming clouds from which it eventually falls back to the ground as rain. By this process, water vapour is naturally transformed into liquid water even when the relative humidity at ground level is below 100%. This prevents the air from reaching equilibrium, ensuring that water can evaporate indefinitely. The constant flow of the water cycle keeps the atmosphere’s water forever out of balance.
How To Be Cool
The water cycle is actually not so different from the cycles of modern artificial refrigerators and air conditioners. In both cases, water evaporates in one region that induces cooling and then condenses back to liquid in another region where it results in heating. The trick is that the heating occurs in a different place from the cooling where we can happily tolerate it. In the case of the water cycle, the cooling evaporation of water happens close to ground level (where we can enjoy it), while the warming condensation occurs in the upper troposphere (where we need not worry about it). Similarly, in a refrigerator or air conditioner, evaporation of a refrigerant fluid happens inside the area that is to be cooled while the warming condensation occurs outside.
In both cases, this is done by a source of energy that creates different conditions in the different regions of the cycle. The Sun heats the air at ground level more than in the upper troposphere, while the electric motor of a refrigerator or air conditioner compresses the fluid to a higher pressure in the condensing region than in the evaporating region. This differentiation prevents the system from reaching equilibrium, instead allowing for a cyclical flow. This is a core design principle of useful systems; equilibrium systems by their nature do not do anything helpful, but systems in dynamic flow undergo changes that we can put to use.
At large scales, the forces of nature are largely beyond our control. The Sun heats the Earth. Water evaporates, rises and forms clouds. The clouds burst and the rain falls. We cannot make a hot day cold, or a wet day dry. But the inexorable flows of the world hide surprising tools for the resourceful. The violent energy of a gale can give us the power to heat a home, if we put a wind turbine in its path. The dangerous penetrating power of x-ray radiation can help us to heal an injury, if we control it in a doctor’s office. The limitations of quantum measurements can allow us to securely communicate, if we use them to protect encryption keys. And even the heat of the sun can cool us, if we moisten our skin a little. Nature can do amazing things — if you don’t mind working up a sweat.
Specifically, the Herald wrote that “iced punch… must have been what the ancients called nectar.”
While later writers remained mainly complementary of ice, they also raised some concerns that it might be too desirable to resist! The Adelaide Observer wrote in 1861 that “It is necessary to guard against the abuse of this luxury. What with its novelty and its tempting coolness, the appearance of ice as a regular commodity in the confectioners’ shops of Adelaide has led in some cases to the consumption of the article not conducive to health. To eat sixpennyworth of ice, as sweetmeats are eaten, is a most ridiculous mistake with regard to the use of the article, and is moreover a very dangerous feat at this time of year.”
An exception to this is in the night and early morning, when the temperature can fall so low that it can no longer hold all of the moisture it contained during the day. This can lead to condensation occuring in the form of dew, which is why the ground often feels wet in the morning even when there has not been any rain.
I have two beginner's questions:
1. Why did they import ice from Boston if antarctica is relatively close-by?
2. if you write that heat is emitted or absorbed, that is in the form of (infrared) radiation, right?
Very cool article