Thermochemistry
Distinguish heat, temperature, system, surroundings, endothermic, and exothermic processes. Calculate heat using q = mcDeltaT and interpret energy diagrams.
A lesson of Chemistry, in Energy, States of Matter, and Solutions, about 35 minutes.
By the end of this lesson, you will be able to say which way energy flows in a reaction, read the sign of $\Delta H$ correctly, calculate heat from a temperature change, and use Hess's law.
Bonds Cost Energy to Break and Release It to Form
That is the whole mechanism, and it is the reverse of most people's intuition.
$$\textbf{breaking a bond ABSORBS energy}$$
$$\textbf{forming a bond RELEASES energy}$$
A bond exists because the arrangement is stable. Pulling it apart means working against that stability, which costs energy. Letting atoms fall together into a bond releases it.
Every reaction does both, and the net result is the difference:
$$\Delta H = (\text{energy to break reactant bonds}) - (\text{energy released forming product bonds})$$
Exothermic means the products' bonds released more than the reactants' cost. Energy leaves the system, the surroundings get warmer, and $\Delta H$ is negative.
Endothermic means breaking cost more than forming released. Energy enters the system, the surroundings get colder, and $\Delta H$ is positive.
The sign is written from the SYSTEM's point of view, which is why exothermic is negative: the reaction lost that energy. A beaker that feels hot is losing energy to your hand, and the negative sign is describing the reaction rather than the beaker.
Sort It: Exothermic or Endothermic
Freezing is exothermic, which surprises people because the result is cold. The water releases energy as its molecules lock into place; a freezer works by carrying that released energy away.
Reading an Energy Diagram
Reactants on the left, products on the right, energy up the side.
$$\textbf{exothermic}: \text{products BELOW reactants}, ; \Delta H < 0$$
$$\textbf{endothermic}: \text{products ABOVE reactants}, ; \Delta H > 0$$
Between them sits a hump. Its top is the activation energy $E_a$: the energy needed to get the reaction started, because bonds must break before new ones can form.
$E_a$ and $\Delta H$ are independent. A strongly exothermic reaction can have an enormous activation energy and sit there doing nothing. Gasoline in air is thermodynamically desperate to burn and needs a spark, and a mixture of hydrogen and oxygen can sit unchanged for years until one flame arrives.
That gap is why the world is not on fire. How far downhill a reaction goes is a different question from how easily it starts, and the kinetics lesson is about the second one.