Thermal Energy, Part 1: What is heat?

This article describes heat and temperature for anyone wanting a deeper understanding of thermoregulation. Doing so requires a quick explainer of some fundamental principles of physics. We’ll tackle those first before turning to the practical matter of why it matters.

What is heat?

At the most basic, heat is energy. Often called thermal energy, heat is a type of kinetic energy (kinetic energy is the energy of motion) applied at the molecular level. Heat has been described as “our macroscopic perception of moving particles- the faster they are moving, the hotter a substance feels to us”.

You’re already familiar with kinetic energy, even if you never think about it, since it is inherent to anything moving. When you drive, you experience kinetic energy as vehicle motion. To stop, you apply a counteracting force, friction from the brakes, converting nearly all the car’s kinetic energy to thermal energy. This is why overused brakes will heat up and smoke (a small amount of energy is also converted to sound, like the squeak you hear when you need to change brake pads).

The amount of a car’s kinetic energy depends on two factors- the mass (size and weight) of the car and the velocity (speed) it moves. Of these two factors, velocity has a much greater effect on the amount of energy, since the kinetic energy equation squares the velocity component (KE = ½mv²). We used a car example, but the same principle applies to anything with mass and motion – literally everything in the universe at least as big as an atom.[1]

We must consider matter much smaller than a car to understand heat. For energy in atoms and molecules, the velocity component comes from vibration (in solid objects) or vibration and movement (in liquids and gasses). Any atom with a temperature above absolute zero vibrates, even if only slightly (absolute zero is -459.67^oF, or 0^oK using the scientifically-preferred Kelvins temperature scale). As for mass, even infinitesimally small atoms have a tiny amount of it.

The combination of mass and velocity mean energy is inherent in all atoms and molecules, even if only a very small amount. According to the third law of thermodynamics, reaching absolute zero is impossible, so we know every atom and molecule in the universe has a temperature at least a bit above -459.67^oF!

Measuring Heat

Heat is commonly measured in two units, the joule and the calorie. Like miles and kilometers both measure distance, just according to a different standard, joules and calories both measure energy (one calorie equals a little more than four joules).

It’s no coincidence we use calories to measure food energy; the original method to determine the number of calories in food involved burning it and measuring the heat released. Humans convert food energy, stored as chemical bonds in fat, protein, and carbohydrates, to useable energy (much of it heat) needed for life by also “burning calories”.[2]

Temperature

Heat is energy, and temperature is how that energy is measured. Temperature measures the average thermal energy of a substance; it is a “rough index of average atomic velocities” in a substance. That “average” part is important in thermodynamics. To show why, let’s consider an ICEPLATE® full of just-melted 32^oF liquid water.

A single 1.5-liter ICEPLATE® has about fifty septillion individual water molecules (that’s 5 followed by 25 zeros!). Some of these molecules move faster than average (a few much faster), and some slower, but on average the thermal energy of all 50-septillion molecules works out to a temperature of 32^oF.

The different energies in each individual water molecule are the result of random collisions between molecules as they move around. The textbook example uses a pool table and billiard balls. A head-on collision between two speeding balls will slow both down, while two balls quickly knocking into another from the left sends that third ball speeding off to the right. The same is true, in principle, for the random movement of water molecules. It’s impossible to measure any single H₂O molecule, but with 50-septillions of them banging around in the ICEPLATE®, we get a good idea of the average.   

The Physics of Sweat Evaporation

The concept of temperature as the average thermal energy of a substance is fundamental to understanding how sweating cools you down. In hot conditions, sweat evaporation is the only natural method humans have for dispelling metabolic heat. This is why Qore Performance products that promote airflow allow you to work harder, longer, and safer in hot conditions.

To understand why, consider a single drop of sweat.

Maybe you’ve been working hard on a hot day, enough to raise your body core temperature to about 100^oF. The H₂O molecules in your sweat drop will also start with an average temperature of 100^oF. However, some H₂O molecules have higher than average energy. Your overheated body pumps more thermal energy into the sweat drop by circulating hot blood near skin surface. The highest energy H₂O molecules soon have enough energy (i.e., speed)[3] to phase change into vapor and fly off, i.e., evaporate.

Since only high energy molecules evaporate, the average temperature of the sweat drop declines a bit. That’s great for thermoregulation, since our overheated bodies keep transferring more energy to the drop of sweat, raising its temperature.

The process is continuous. The highest energy H₂O molecules evaporate away, taking energy with them and decreasing the average temperature of the sweat drop, while the body transfers more energy into the sweat drop, increasing its temperature.

Eventually, your single drop of sweat dries up, completely evaporated away. If core temperature remains elevated (and you are hydrated), you keep sweating. Energy transfer continues from your core, through blood circulation, to sweat evaporation and, ultimately, to the environment. Through sweating, we avoid overheating if the rate of heat transfer through evaporation is greater than, or equal to, the rate of metabolic heat production in the body.

This process assumes sweat is free to evaporate. If sweat can’t evaporate, those high-energy H₂O molecules aren’t removed and no energy transfer occurs. High humidity reduces sweat evaporation since the air is already nearing saturation. Often, safety gear and clothing can also restrict sweat evaporation.

When working outdoors in high humidity, there is little we can do to promote sweat evaporation except to increase airflow. Qore Performance products like ICEVENTS® help by promoting airflow, enabling sweat to evaporate and provide cooling.

Putting it all together – the “So What”?

So far, this article has been abstract. To demonstrate why this matters, consider the water in a Qore Performance ICEPLATE® Curve. In this example, we’ll assuming we just pulled an ICEPLATE® out of the cooler and geared up. The water molecules (all fifty septillion of them) have an average temperature of 0^oF (remember, temperature measures thermal energy). Trapped in the crystal structure of ice, each H₂O molecule vibrates but isn’t free to move around.

Above: The Qore Performance ICEPLATE® Curve, a 50 fl oz / 1.5L water bladder that can be filled with frozen or heated water against the body for thermoregulation benefits in extreme temperatures.

The second law of thermodynamics states that heat flows from hot to cold,[4] so after strapping on the ICEPLATE®, heat immediately begins flowing from body to ICEPLATE®. This happens through a process called conductive heat transfer. The atoms and molecules in your body, being around 98.6^oF, vibrate more energetically than the 0^oF ice in the ICEPLATE®. This energetic (i.e., temperature) difference between body and ice results in energy transfer into the ICEPLATE®, causing frozen H₂O molecules to begin vibrating faster. Eventually, the temperature of the ice reach 32^oF.

By now, raising the temperature of the ice from 0^oF to 32^oF has transferred 13,322 calories into the ice (recall calories = a unit of energy). That’s about 13kcals – the energy equivalent of about four individual M&Ms. Not much!

Fortunately, the next stage of the process draws much more thermal energy. Each H₂O molecules now has a temperature of 32^oF but the molecules remain locked in their ice crystal structure. Since we are still wearing the ICEPLATE®, energy is still flowing from our body (98.6^oF) to the ice (now 32^oF but still frozen). Now, energy absorbed into the ICEPLATE® works to break molecular bonds holding H₂O molecules frozen together.

Releasing H₂O molecules from their frozen bonds requires a huge amount of energy. Completely melting all water in an ICEPLATE® takes more than 119,741 calories (about 120 kcals, over 800% more energy than it took to warm the ice from 0^oF to 32^oF).

Immediately after melting, H₂O molecules still have a temperature of 32^oF but are no longer bonded. If we keep wearing the ICEPLATE®, we’ll pump even more heat into the now-liquid system. Freed from the bonds that held them as solid ice, additional thermal energy makes H₂O molecules both vibrate and move faster. Raising the temperature of liquid water, with molecules free to move around, requires more energy than raising the temperature of solid water but far less energy than melting ice.

If we keep wearing our plate until the water temperature reaches a refreshing (but no longer cold) 50^oF, the ICEPLATE® draws in another 15,000 calories (15kcals; the energy equivalent of another 5 M&Ms).

All told, from 0^oF, through the melting process, to 50^oF water requires the ICEPLATE® to absorb about 148,000 calories of thermal energy. This value approximates the typical thermal energy absorbed when using single ICEPLATE® products like the ICEPLATE® Backpack Gen 3. Products that can hold two ICEPLATEs, like the ICEPLATE EXO® Gen 3 for military or the ICEPLATE® SLK Gen 3 for business and safety, double that amount of thermal energy absorption to approximately 296,000 calories!

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About the author: Dr. Erik Patton holds a PhD from Duke University where he conducted research on the challenges rising temperatures pose for military training. An Army veteran, Erik has served in a variety of extreme climates ranging from deserts in the U.S. Southwest and Middle East (120^oF) to Arctic conditions in central Alaska (-42^oF).