Exothermic: $\Delta H < 0$, heat is released. Endothermic: $\Delta H > 0$, heat is absorbed.

Bond Enthalpies: Uses bond breaking and forming energies. Enthalpies of Formation: Uses standard enthalpies of formation of products and reactants.

A process that releases heat to the surroundings; $\Delta H$ is negative.

A process that absorbs heat from the surroundings; $\Delta H$ is positive.

Energy is conserved; it cannot be created or destroyed, only transferred or converted.

The state where heat flows until two objects reach the same temperature.

The energy needed to raise 1 gram of a substance by 1 degree Celsius (J/g°C).

The heat change at constant pressure; indicates whether a reaction is exothermic or endothermic.

The heat change when one mole of a compound is formed from its elements in their standard states.

A property whose value does not depend on the path taken to reach a specific value.

Measure the initial temperature of reactants and calorimeter, conduct the reaction inside the calorimeter, measure the final temperature, and calculate the heat released or absorbed using $q = mc\Delta T$.

Sum of bond energies broken minus the sum of bond energies formed: $\Delta H = \Sigma \text{Bonds Broken} - \Sigma \text{Bonds Formed}$.

Sum of enthalpies of formation of products minus the sum of enthalpies of formation of reactants: $\Delta H = \Sigma \Delta H_f^\circ \text{(Products)} - \Sigma \Delta H_f^\circ \text{(Reactants)}$.

If a reaction can be expressed as a series of steps, the enthalpy change for the overall reaction is the sum of the enthalpy changes for each step. Manipulate individual reactions (reversing or multiplying) to match the target reaction.