Enthalpy (ΔH) - MCAT General Chemistry

Enthalpy (ΔH)

10 min read Updated Mar 26, 2026

Every chemical reaction involves energy. Some reactions release energy and feel hot to the touch. Others absorb energy and feel cold. Enthalpy is the thermodynamic quantity that tracks this heat flow at constant pressure - which is how most reactions in your body and in the lab actually occur.

What Is Enthalpy?

Enthalpy (H) is defined as:

In practice, you almost never calculate H directly. Instead, you work with the change in enthalpy:

Δ $H_{\text{rxn}}$ = H(products) - H(reactants)

Exothermic vs. Endothermic

Property	Exothermic	Endothermic
ΔH sign	Negative (ΔH < 0)	Positive (ΔH > 0)
Energy flow	System releases heat to surroundings	System absorbs heat from surroundings
Surroundings feel	Warmer	Cooler
Products vs. reactants	Products at LOWER energy	Products at HIGHER energy
Example	Combustion of methane	Photosynthesis

Enthalpy Diagrams

An enthalpy diagram places reactants and products on a vertical energy axis. For an exothermic reaction, the products sit lower than the reactants - energy was released going “downhill.” For an endothermic reaction, the products sit higher - energy was absorbed going “uphill.”

The vertical distance between reactants and products is |ΔH|. The direction tells you the sign.

Energy diagram showing reactants X at lower energy, products Y at higher energy, with activation energy Ea labeled for both forward and reverse reactions, and delta-H showing the overall enthalpy change. A dashed red line shows the lower activation energy pathway with a catalyst. — An enthalpy (energy) diagram for an endothermic reaction. Reactants (X) sit lower than products (Y), and the difference is ΔH. The peak represents the activation energy (Ea). A catalyst (dashed red line) lowers Ea but does not change ΔH. *Credit: Wikimedia Commons, CC BY-SA 3.0*

Enthalpy and the PV Term

The full definition H = U + PV includes a pressure-volume term. For reactions involving only solids and liquids, volume changes are tiny, so ΔH is almost equal to ΔU. But for reactions involving gases, the PV term matters because gases expand or compress significantly.

For an ideal gas at constant temperature: PV = nRT

So ΔH = ΔU + Δ(nRT) = ΔU + ΔnRT

where Δn = moles of gaseous products - moles of gaseous reactants. If a reaction produces more moles of gas than it consumes (Δn > 0), ΔH > ΔU. If it produces fewer moles of gas (Δn < 0), ΔH < ΔU.

Reversibility of Enthalpy

Because enthalpy is a state function, the enthalpy change for the reverse reaction has the same magnitude but opposite sign:

Forward: H₂(g) + $\frac{1}{2}$ O₂(g) → H₂O(l), ΔH = -285.8 kJ
Reverse: H₂O(l) → H₂(g) + $\frac{1}{2}$ O₂(g), ΔH = +285.8 kJ

If it costs 285.8 kJ to decompose water, then forming water releases exactly 285.8 kJ. This principle is essential for Hess’s law calculations.

Stoichiometry and Enthalpy

ΔH values are reported per mole of reaction as written. If you double all the coefficients, you double ΔH:

H₂(g) + $\frac{1}{2}$ O₂(g) → H₂O(l), ΔH = -285.8 kJ
2 H₂(g) + O₂(g) → 2 H₂O(l), ΔH = -571.6 kJ

Always check the stoichiometry when using ΔH values in calculations.

If dissolving ammonium nitrate in water has ΔH = +25.7 kJ/mol, what happens to the temperature of the water?

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The water temperature decreases. A positive ΔH means the reaction is endothermic - it absorbs heat from the surroundings (the water). As heat is pulled from the water into the dissolving process, the water cools down. This is exactly how instant cold packs work.

For the reaction N₂(g) + 3 H₂(g) → 2 NH₃(g), ΔH = -92.4 kJ. What is ΔH for the decomposition of 1 mol of NH₃ into N₂ and H₂?

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ΔH = +46.2 kJ. The reverse reaction has the opposite sign: +92.4 kJ for decomposing 2 mol NH₃. Since we only want 1 mol, divide by 2: 92.4 / 2 = +46.2 kJ.