Specific heat capacity is the reason a metal spoon gets scalding hot in seconds while the soup it sits in stays lukewarm for minutes. Defined as the energy (in joules) needed to raise one gram of a substance by one degree Celsius (or one kelvin), it tells you how much thermal energy a material can absorb before its temperature noticeably changes. Water has a famously high specific heat capacity of 4.184 J/(g·°C) — one of the highest of any common substance. By comparison, iron's specific heat is only 0.449 J/(g·°C), meaning the same amount of energy that raises water's temperature by 1°C would raise iron's temperature by over 9°C.
Water's extraordinary specific heat has planetary consequences. Oceans absorb enormous amounts of solar energy during the day and release it slowly at night, moderating coastal climates. This is why San Francisco (next to the Pacific) has mild temperatures year-round, while Sacramento, just 140 km inland, swings from freezing winters to scorching summers. The same property makes water an excellent coolant in car engines, nuclear reactors, and industrial processes. Your body exploits it too — since you are about 60% water, your internal temperature stays remarkably stable despite large fluctuations in your environment and metabolic heat production.
At the atomic level, specific heat depends on how many ways a substance can store energy. Monatomic gases like helium can only store energy as kinetic motion (translation), giving them low specific heats. Complex molecules like water can also rotate and vibrate in multiple modes, soaking up much more energy before their temperature rises. In solids, the Dulong-Petit law predicts that most metallic elements have a molar heat capacity close to 25 J/(mol·K) — roughly 3R, where R is the gas constant — because each atom vibrates in three dimensions. This law works well for heavy elements at room temperature but breaks down for light elements like diamond (carbon), whose strong bonds and low mass require quantum mechanics to explain.