---
title: "Enthalpy of Solution — AP Chem Definition & Exam Guide"
description: "Enthalpy of solution is the heat absorbed or released when one mole of solute dissolves. It links calorimetry, IMFs, and free energy of dissolution on the AP exam."
canonical: "https://fiveable.me/ap-chem/key-terms/enthalpy-of-solution"
type: "key-term"
subject: "AP Chemistry"
unit: "Unit 7"
---

# Enthalpy of Solution — AP Chem Definition & Exam Guide

## Definition

Enthalpy of solution (ΔH°soln) is the enthalpy change when one mole of a solute dissolves in a solvent to form a dilute solution. It can be positive (endothermic, like KCl in water) or negative (exothermic), and it is one piece of whether dissolution is thermodynamically favorable.

## What It Is

Enthalpy of solution, written ΔH°soln, is the [heat absorbed](/ap-chem/unit-6/energy-phase-changes/study-guide/kAXAzHrD24XL6LdpMFHh "fv-autolink") or released when exactly one [mole](/ap-chem/key-terms/mole "fv-autolink") of a solute dissolves in a solvent to make a dilute solution. Think of dissolving as an energy trade with three parts. First, solute particles have to be pulled apart from each other (that costs energy). Second, solvent molecules have to be pushed apart to make room (that also costs energy). Third, new solute-solvent attractions form (that releases energy). Add up the costs and the payoff and you get ΔH°soln.

The sign tells you which side won. If the new solute-solvent attractions release more energy than it took to break the old [interactions](/ap-chem/unit-3/representations-solutions/study-guide/O4uZStuqpe603GRuztjG "fv-autolink"), ΔH°soln is negative and the solution warms up. If breaking interactions cost more, ΔH°soln is positive and the solution gets colder. KCl dissolving in water is the classic endothermic example, which is why instant cold packs work. The key AP insight is that a positive ΔH°soln does not stop a salt from dissolving, because entropy gets a vote too.

## Why It Matters

This term lives in Topic 7.14, Free Energy of Dissolution, which sits at the end of the solubility equilibrium topics in [Unit 7](/ap-chem/unit-7 "fv-autolink"). The whole point of 7.14 is that dissolution is a process governed by thermodynamics, so its favorability depends on both ΔH°soln and the entropy change of dissolving, combined through ΔG = ΔH − TΔS. Enthalpy of solution is the bridge concept that ties together three things you learned separately. It connects [intermolecular forces](/ap-chem/key-terms/intermolecular-forces "fv-autolink") (why things dissolve), calorimetry (how you measure the heat), and free energy (whether dissolution is thermodynamically favored and how that relates to Ksp). The College Board loves this term precisely because it forces you to use multiple units at once.

## Connections

### Calorimetry and q = mcΔT (Unit 6)

Enthalpy of solution is measured with coffee-cup calorimetry. You dissolve a known mass of [solute](/ap-chem/key-terms/solute "fv-autolink"), track the temperature change of the water, and use q = mcΔT to find the heat, then divide by moles of solute. A temperature drop means the dissolution pulled heat from the water, so ΔH°soln is positive.

### Intermolecular Forces and Solutions (Unit 3)

ΔH°soln is basically an IMF accounting problem. You break solute-solute and solvent-solvent attractions and form solute-solvent attractions. '[Like dissolves like](/ap-chem/key-terms/like-dissolves-like "fv-autolink")' works because similar IMFs make the energy payoff from new solute-solvent interactions big enough to cover the cost of breaking the old ones.

### Free Energy and Entropy, ΔG = ΔH − TΔS (Topic 7.14 / Unit 7)

ΔH°soln alone does not decide whether a [salt](/ap-chem/key-terms/salt "fv-autolink") dissolves. KCl dissolves even though its ΔH°soln is positive because dissolving scatters ions through the solution and increases entropy. When TΔS outweighs ΔH, ΔG is negative and dissolution is thermodynamically favored.

### Solubility Equilibria and Ksp (Unit 7)

Topic 7.14 connects the thermodynamics of dissolving to the equilibrium picture. A salt whose dissolution has a negative ΔG tends to have a larger Ksp, and an endothermic ΔH°soln predicts that solubility increases as temperature rises (a Le Châtelier-style argument).

## On the AP Exam

This concept shows up in both multiple choice and free response, almost always as a multi-skill problem rather than a definition question. The 2026 long FRQ opened with a calorimetry experiment on KCl, the textbook endothermic dissolution, and that setup is typical. You should be ready to do four things. One, calculate ΔH°soln from calorimetry data using q = mcΔT and moles of solute (watch signs: temperature drop means positive ΔH°soln). Two, explain the sign of ΔH°soln in terms of breaking and forming particle-level interactions, since particle-level reasoning is what earns the explanation points. Three, explain why an endothermic dissolution can still occur, using entropy and ΔG = ΔH − TΔS. Four, predict how temperature affects solubility from the sign of ΔH°soln. MCQ stems often give you a temperature change and ask for the sign of ΔH, or describe a salt with positive ΔH°soln and ask why it dissolves anyway.

## enthalpy of solution vs Enthalpy of fusion (melting)

Both involve a solid 'turning into' mobile particles, but they are different processes. Melting is a phase change of a pure substance, where only solute-solute (or molecule-molecule) interactions are disrupted, so enthalpy of fusion is always positive. Dissolution involves a second substance, the solvent, so you break solute-solute AND solvent-solvent interactions but also form new solute-solvent attractions. That extra energy payoff is why ΔH°soln can be either positive or negative, while melting is always endothermic.

## Key Takeaways

- Enthalpy of solution (ΔH°soln) is the enthalpy change when one mole of solute dissolves in a solvent to form a dilute solution.
- Its sign comes from an energy trade. Breaking solute-solute and solvent-solvent interactions costs energy, while forming solute-solvent interactions releases energy.
- Dissolution can be endothermic or exothermic. KCl dissolving in water is endothermic, which is why the solution gets colder and why cold packs work.
- A positive ΔH°soln does not prevent dissolving. If the entropy increase makes TΔS larger than ΔH, then ΔG is negative and dissolution is thermodynamically favored.
- You measure ΔH°soln with coffee-cup calorimetry. Use q = mcΔT on the water, flip the sign for the process, and divide by moles of solute.
- If ΔH°soln is positive, raising the temperature generally increases the salt's solubility, which connects enthalpy of solution back to Ksp and equilibrium shifts.

## FAQs

### What is enthalpy of solution in AP Chem?

It is the enthalpy change when one mole of a solute dissolves in a solvent to form a dilute solution, written ΔH°soln. It appears in Topic 7.14, Free Energy of Dissolution, where it combines with entropy to determine whether dissolving is thermodynamically favorable.

### Is dissolving always exothermic?

No. Dissolution can go either way. KCl and NH4NO3 dissolve endothermically (the solution cools, which is how instant cold packs work), while salts like NaOH and CaCl2 dissolve exothermically. The sign depends on whether forming new solute-solvent attractions releases more energy than breaking the original interactions cost.

### How is enthalpy of solution different from enthalpy of fusion (melting)?

Melting is a phase change of one pure substance, so it only disrupts the particles' own interactions and is always endothermic. Dissolution involves a solvent, so new solute-solvent attractions form and release energy, which is why ΔH°soln can be positive or negative.

### If ΔH of solution is positive, why does the salt still dissolve?

Because entropy drives it. Dissolving an ionic solid scatters ions throughout the solution, so ΔS is usually positive. In ΔG = ΔH − TΔS, a large enough TΔS term makes ΔG negative even when ΔH is positive, so dissolution is still thermodynamically favored.

### How do you calculate enthalpy of solution from a calorimetry experiment?

Dissolve a measured mass of solute in water, record the temperature change, and compute q = mcΔT for the water. Then flip the sign (heat lost by the water is heat gained by the dissolution) and divide by moles of solute to get ΔH°soln in kJ/mol. The 2026 long FRQ used exactly this setup with KCl.

## Related Study Guides

- [7.14 Free Energy of Dissolution](/ap-chem/unit-7/free-energy-dissolution/study-guide/tsCK6wFJQBajLTmMNv44)

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