Enthalpy of reaction (ΔH or ΔHrxn) is the heat energy absorbed or released by a chemical reaction at constant pressure; on the AP Chem exam, a negative ΔH means exothermic, a positive ΔH means endothermic, and ΔH for a multistep process equals the sum of the ΔH values of its steps (Hess's Law).
Enthalpy of reaction (ΔH) is the total heat energy transferred when a reaction happens at constant pressure. If ΔH is negative, the reaction releases heat to the surroundings (exothermic). If ΔH is positive, the reaction absorbs heat (endothermic). The sign convention is the whole game here, so get it locked in early.
The reason this term shows up everywhere in Unit 6 is that ΔH is a state function. It only depends on where you start and where you end, not the path you take. That's the entire logic behind Hess's Law (Topic 6.9). Because energy is conserved (the first law of thermodynamics), you can break any process into steps, and the overall ΔH equals the sum of the ΔH values of those steps (EK 6.9.B.1). Think of it like elevation on a hike. Whether you take the steep trail or the winding one, the change in altitude from trailhead to summit is the same.
Enthalpy of reaction is the central quantity in Unit 6 (Thermochemistry) and the backbone of Topic 6.9, Hess's Law. Learning objective 6.9.A asks you to represent a process as a sequence of steps, and 6.9.B asks you to explain why the overall enthalpy equals the sum of the step enthalpies. Both objectives are really asking the same thing in different clothes. Can you do algebra with ΔH values and justify it with conservation of energy?
It also reaches into Unit 7. Topic 7.6 has you manipulate equilibrium constants the same way you manipulate thermochemical equations under Hess's Law. Reverse a reaction and ΔH flips sign (while K gets inverted). Multiply coefficients by c and ΔH gets multiplied by c (while K gets raised to the power c). Add reactions and you add the ΔH values (while you multiply the K values). If you can manipulate one, you can manipulate the other.
Keep studying AP Chemistry Unit 7
Hess's Law (Unit 6)
Hess's Law is enthalpy of reaction in action. Because ΔH is a state function, the ΔH of an overall reaction equals the sum of the ΔH values of any set of steps that gets you from reactants to products. When you reverse a step, flip the sign of its ΔH; when you scale a step, scale its ΔH by the same factor.
Properties of the Equilibrium Constant (Unit 7)
Topic 7.6 uses the exact same equation-manipulation rules, just with K instead of ΔH. Adding reactions means adding ΔH values but multiplying K values, and reversing a reaction flips the sign of ΔH but inverts K. Recognizing this parallel makes both topics easier to study.
Exothermic and Endothermic Reactions (Unit 6)
These labels are just the sign of ΔH translated into words. ΔH < 0 means exothermic (heat leaves the system, surroundings warm up), and ΔH > 0 means endothermic (heat enters the system, surroundings cool down). Calorimetry FRQs love asking you to connect a temperature change to the sign of ΔH.
Activation Energy (Unit 5)
On an energy diagram, ΔH is the height difference between reactants and products, while activation energy (Ea) is the hill between them. ΔH tells you thermodynamics (how much heat overall), and Ea tells you kinetics (how fast the reaction can go). A reaction can have a very negative ΔH and still be slow because Ea is huge.
Enthalpy of reaction shows up in two main flavors. First, straight Hess's Law math. A question gives you steps like A → B (ΔH = +50 kJ) and B → C (ΔH = -120 kJ) and asks for ΔH of A → C, which is just the sum (-70 kJ). The twist is usually manipulation: you may need to reverse a step (flip the sign) or multiply a step (scale the ΔH) before adding. Second, experimental design. The 2023 FRQ on CaCO₃(s) + HCl(aq) and practice questions in the same style ask you to plan calorimetry measurements for two easy-to-run reactions, then combine them with Hess's Law to find the ΔH of a reaction you can't measure directly. A classic trick is measuring A(s) + C(aq) → D(aq) and B(s) + C(aq) → D(aq), then subtracting to get A → B. Multiple-choice stems also test the validity condition, which is that the steps must actually add up to the target equation with intermediates canceling. Acid-base FRQs like the 2024 lactic acid problem can fold in ΔH from calorimetry data too, so be ready to compute q = mcΔT and convert it to ΔH per mole.
Enthalpy of reaction (ΔH) is the net energy difference between products and reactants, which tells you whether heat is released or absorbed overall. Activation energy (Ea) is the energy barrier the reaction must climb to get started, which tells you about reaction rate. On an energy diagram, ΔH compares the start and end levels; Ea is the peak in between. Changing the path (like adding a catalyst) lowers Ea but never changes ΔH, because ΔH is a state function.
Enthalpy of reaction (ΔH) is the heat absorbed or released by a reaction at constant pressure, with negative ΔH meaning exothermic and positive ΔH meaning endothermic.
ΔH is a state function, so the overall enthalpy change depends only on reactants and products, not on the pathway taken between them.
Hess's Law says the ΔH of an overall reaction equals the sum of the ΔH values of its individual steps, which follows directly from the first law of thermodynamics.
When you reverse a reaction, flip the sign of ΔH; when you multiply the coefficients by a factor, multiply ΔH by that same factor.
The same manipulation rules apply to equilibrium constants in Topic 7.6, except adding reactions means multiplying K values instead of adding them.
A catalyst lowers activation energy but never changes the enthalpy of reaction, because ΔH only depends on the initial and final states.
Enthalpy of reaction (ΔH) is the total heat energy transferred during a reaction at constant pressure. A negative ΔH means the reaction is exothermic and releases heat; a positive ΔH means it's endothermic and absorbs heat. It's the core quantity in Unit 6 Thermochemistry.
No. A catalyst lowers the activation energy (Ea), which speeds up the reaction, but ΔH stays exactly the same because it's a state function that depends only on the reactants and products, not the pathway.
ΔH is the net energy difference between products and reactants (thermodynamics), while activation energy is the barrier the reaction must overcome to proceed (kinetics, Unit 5). A reaction with a very negative ΔH can still be extremely slow if its Ea is large.
Manipulate the given steps so they add up to your target equation, then add their ΔH values. Reversing a step flips the sign of its ΔH, and multiplying coefficients by a factor multiplies ΔH by that factor. For example, if A → B has ΔH = +50 kJ and B → C has ΔH = -120 kJ, then A → C has ΔH = -70 kJ.
No, but they follow parallel manipulation rules in Topic 7.6. When you add reactions, you add ΔH values but multiply K values; when you reverse a reaction, ΔH changes sign while K gets inverted. The exam rewards knowing both sets of rules side by side.