Energy changes

The law of conservation of energy states that energy can be neither created nor destroyed. Therefore, the total energy of the universe is constant.

When studying transfers of energy, it is important to distinguish between the system and the surroundings. The system is the part of the universe being studied; the surroudings is everything else in the universe. If the system is the contents of the flask, then the surroundings includes the flask, the air, the table, the room, etc.

There are three types of systems: open, closed, and isolated. The name tells us whether matter and energy can move in and out of the system.

Isolated systems don’t exist in nature. We usually work with closed systems. A stoppered flask, for example, is a closed system because matter is trapped inside and no outside matter is let in, but energy can move in and out freely (otherwise, placing the flask over a flame would have no effect on its contents).

Any change in energy in the system is accompanied by an equal and opposite change in the surroundings: $\mathrm{\Delta }{E}_{\text{system}}+\mathrm{\Delta }{E}_{\text{surroundings}}=0\text{.}$

Reactions are classified based on the system–surroundings energy transfer. In an exothermic reaction, energy is given off (released) by the system. In an endothermic reaction, energy is absorbed by the system. The change in temperature of the surroundings indicates the type of reaction. An increase (measured before and after the reaction) indicates that is was exothermic; a decrease, endothermic.

These state changes are endothermic:

• solid to liquid: melting/fusion
• liquid to gas: vaporization/evaporation/boiling
• solid to gas: sublimation

Going the other way is exothermic:

• liquid to solid: freezing/solidification
• gas to liquid: condensation
• gas to solid: sublimation/deposition