What you'll be able to do: Define the system and surroundings for a chemical change and assign the correct sign to heat transferred.
Almost every thermochemistry mistake is a sign mistake, and almost every sign mistake comes from not deciding what the system is. This lesson sets up that vocabulary once so the calculations later stay clean.
These are recommended, not required. You can start this lesson at any time.
The system is the part of the universe you are studying, normally the chemicals reacting. The surroundings are everything else that can exchange energy with it: the solvent, the container, the air. Deciding this first is what fixes every sign that follows.
Energy cannot be created or destroyed, only moved or converted. For a chemical change this means the energy released by breaking and forming bonds has to appear somewhere else, usually as warmer water. Written out, the internal energy change of the system equals the heat added plus the work done on it.
DE = q + w
Signs are written from the point of view of the system. Heat entering the system is positive, so an endothermic change has q > 0 and the surroundings cool. Heat leaving the system is negative, so an exothermic change has q < 0 and the surroundings warm.
Temperature is an intensive property, the average kinetic energy of the particles. Heat is an amount of energy transferred. A bathtub of warm water at 40 C holds far more thermal energy than a cup at 90 C even though the cup is hotter.
Breaking bonds always costs energy and forming bonds always releases it. A reaction is exothermic when the bonds formed in the products are collectively stronger than the bonds broken in the reactants. That single idea will return in the bond enthalpy lesson.
The chemicals or process under study; everything else is the surroundings.
Energy transferred because of a temperature difference, measured in joules.
q is positive when heat flows into the system and negative when heat flows out.
The system releases heat, q(system) < 0, and the surroundings get warmer.
The system absorbs heat, q(system) > 0, and the surroundings get cooler.
DE = q + w
q(system) = -q(surroundings)
A solid dissolves in water and the water temperature falls from 22.0 C to 18.4 C. Is dissolving endothermic or exothermic, and what is the sign of q for the dissolving process?
Try it first: Decide which part is the system before looking at the thermometer.
0 of 4 steps revealed.
Burning 1.00 g of methane in a calorimeter raises the water temperature. State the signs of q(reaction) and q(water), and explain why they differ.
0 of 3 steps revealed.
Why it's wrong: A rising thermometer describes the surroundings, not the system.
Check instead: Water warms means the reaction released heat, so the reaction q is negative.
Why it's wrong: Temperature is intensive; heat depends on how much material there is.
Check instead: Ask whether the quantity would change if you doubled the sample.
Why it's wrong: The water is the measuring device, it is what absorbs the released heat.
Check instead: System equals the reacting species only.
Why it's wrong: It draws heat out of its surroundings, including your hand.
Check instead: Endothermic feels cold to the touch.
No practice questions are available for this topic yet. You can still practice the whole unit.
Energy is conserved: whatever the system loses, the surroundings gain. The system is the reacting chemicals, the surroundings are everything else, usually the water and the calorimeter. Heat flowing into the system is positive, heat flowing out is negative. An exothermic reaction releases heat, so q(system) is negative while the surrounding water warms up. Endothermic is the reverse. Temperature measures average kinetic energy; heat is energy in transit because of a temperature difference. The two are not the same thing.
Chemistry 2e, OpenStax, Rice University, licensed CC BY 4.0. License