Heat Transfer and Thermal Equilibrium

Temperature measures the average kinetic energy of particles. When two objects touch, particles at the boundary collide. Faster particles from the hot object hit slower particles from the cold one, and on average energy passes from fast to slow, never the other way. This keeps happening, billions of collisions per second, until the average kinetic energy on both sides is equal. That shared condition is thermal equilibrium: no more net energy transfer, because neither side has a speed advantage anymore.

Take two metal blocks, one at 80 C and one at 20 C, pushed together. At the contact surface, particles from the hot block slam into particles from the cold block and lose some energy with each collision. The hot block's particles slow down, the cold block's speed up. The effect spreads inward from the boundary. Eventually both blocks sit at the same intermediate temperature. Nothing physical called heat sits inside either block before or after. Heat is just the name for energy moving across that boundary while a temperature difference exists. Once temperatures match, the transfer stops, even though both blocks still contain plenty of internal thermal energy.

This is why a calorimetry problem is really a bookkeeping exercise: whatever energy leaves the hot substance must enter the cold one, since no energy is lost to the surroundings. Setting heat lost equal to heat gained, using q = mc(deltaT) for each substance, lets you solve for one unknown final temperature shared by both. This same q = mc(deltaT) relationship reappears throughout the unit, in heating curves and in calorimeter problems, because it is the basic tool for tracking energy moving as heat.