Reversing a chemical equation reverses the direction of energy flow between system and surroundings.

A system is drawn inside its surroundings with labeled arrows for heat and work crossing the boundary. The diagram emphasizes that energy flow direction determines the sign convention (inward heat/work increases the system’s energy; outward decreases it), which is the same logic used when interpreting the sign of $\Delta H$ for reactions written forward vs. reverse. Source

On the AP Exam, you must connect the written reaction direction to the sign of $\Delta H$ and interpret what that means physically.

Core idea: reversing a reaction reverses heat flow

A chemical reaction can be written in the forward direction or the reverse direction. The enthalpy change reported for the reaction is tied to the reaction as written.

• If the forward reaction releases heat to the surroundings, the reverse reaction must absorb that same amount of heat from the surroundings.

• If the forward reaction absorbs heat, the reverse reaction must release that same amount of heat.

• Therefore, the numerical value stays the same, but the sign flips.

What “same magnitude, opposite sign” means

“Same magnitude” means the size of the energy change is unchanged (e.g., 125 kJ stays 125 kJ). “Opposite sign” means positive becomes negative, and negative becomes positive.

Enthalpy change is a property of the initial and final states

This rule works because enthalpy depends on the starting and ending states of the system. Reversing the reaction swaps those states, so the energy difference changes direction.

Because the sign of $\Delta H$ communicates the direction of heat transfer, reversing the chemical equation must reverse the sign to stay consistent with energy conservation.

Thermochemical equations: the reaction and $\Delta H$ are paired

A thermochemical equation is a balanced chemical equation written together with its $\Delta H$ value. The key rule is that the written equation and the stated $\Delta H$ form a matched set.

• Forward equation written with $\Delta H < 0$ describes an exothermic direction.

• Reverse equation written with $\Delta H > 0$ describes an endothermic direction.

• You must reverse the equation exactly (reactants/products swap); then flip the sign of $\Delta H$.

The sign flip rule (how to write it)

Use a compact statement to keep direction and sign consistent.

This relationship applies whether the process is described per reaction as written or per mole of a specified substance, as long as you are truly reversing the same reaction.

Common AP pitfalls when reversing reactions

Forgetting that $\Delta H$ is attached to the equation direction

• If you flip reactants and products but keep the same $\Delta H$ sign, you are describing the wrong heat flow.

• If you keep the equation the same but flip the sign, you are also describing a different process than the one written.

Losing track of “system” vs. “surroundings”

• $\Delta H < 0$: heat leaves the system (surroundings warm).

Endothermic and exothermic processes are shown as system–surroundings energy transfers, with arrows indicating the direction of heat flow. The figure directly ties the sign of $\Delta H$ to the physical observation: surroundings cool when the system absorbs heat ($\Delta H>0$) and warm when the system releases heat ($\Delta H<0$). Source

• $\Delta H > 0$: heat enters the system (surroundings cool).

When you reverse the reaction, which side “gains” chemical potential energy versus “loses” it is reversed, so the heat direction must reverse too.

Not keeping physical states consistent

Reversal assumes the same reactants and products in the same physical states (s, l, g, aq). If states change, it is not the exact reverse process, and the enthalpy change might not be a simple sign flip.

Confusing reversal with other equation changes

For this rule, only the direction changes. The coefficients remain exactly what they were in the original equation; only reactants and products swap sides.