Your body knows that internal organs are more vital than, say, fingers and toes. In cold weather, it wouldn’t be very efficient to pump warm blood to the outside of your body, because then the heat could easily radiate away! Conserving warmth in your core is a great survival tactic, but it’s uncomfortable for the body parts that aren’t receiving as much circulation. Luckily, science has a solution! We’ve seen it before. We’ll see it again. It’s one of the most wildly useful concepts in chemistry! Let’s take another look at exothermic reactions.
Reusable hand warmers are a great everyday example of exothermic behavior. These handy pocket heaters rely on liquid sodium acetate to produce warmth. The liquid is supersaturated– it’s on the verge of crystallizing. When we trigger a crystallization reaction, the sodium acetate changes from liquid (higher energy) to a crunchy near-solid (lower energy). The leftover energy has to go somewhere. It’s released as heat!
The rectangular mark you see here is invisible to the naked eye. The infrared camera, however, “sees” the warm area where the hand warmer contacted the skin!
Sodium acetate releases heat energy as it solidifies. To return a hand warmer to its original, liquid state, all we have to do is put the energy back. Submerge the used hand warmer in boiling water, and the added heat reverses the crystallization reaction! Energy-consuming reactions like this one are called endothermic. After boiling, the packet is back in its original higher-energy state, ready to start the process all over again the next time cold weather strikes.