The free energy of dissolution, , helps explain whether a salt dissolves and how much dissolves under given conditions. It comes from three changes happening at once: breaking apart the solid, reorganizing the solvent, and forming solute-solvent interactions. For AP Chemistry, connect those particle-level changes to enthalpy, entropy, and the sign of .
Why This Matters for the AP Chemistry Exam
This topic asks you to connect particle-level events to a macroscopic result: does the salt dissolve, and how does temperature change that? On the AP Chemistry exam, you are expected to explain the relationship between a salt's solubility and the enthalpy and entropy changes during dissolution. That means justifying claims with reasoning, not just stating that ΔG° = ΔH° - TΔS°.
A big skill here is explaining how well a simple model describes real behavior. Dissolution is a strong example because the three contributions to free energy often nearly cancel, so a clean prediction is hard. Being able to discuss why a model has limits is exactly the kind of reasoning that shows up in free-response explanations.
Key Takeaways
- ΔG° of dissolution reflects three factors: breaking the interactions holding the solid together, reorganizing solvent around the dissolved species, and the dissolved species interacting with the solvent.
- Each factor has both an enthalpy and an entropy piece, and you can estimate the sign and rough size of each.
- These contributions often partly cancel, so predicting the total free energy of dissolution is genuinely hard.
- Use ΔG° = ΔH° - TΔS° to reason about temperature: whether dissolution becomes more or less favored as T rises depends on the signs of ΔH° and ΔS°.
- A negative ΔG° means dissolution is thermodynamically favored under those conditions.
- Solubility is the macroscopic outcome of these particulate-level energy and disorder changes.
The Thermodynamics of Dissolution
Dissolution is not a salt just "disappearing" in water. It is a process with energy and disorder changes you can analyze step by step.
Three Contributions to Free Energy
When a salt dissolves, three things happen at once:
Breaking the solid apart
- The interactions holding the solid together must break.
- This step takes energy (endothermic enthalpy contribution).
- It tends to increase disorder as ordered particles are freed.
Reorganizing the solvent
- Water molecules rearrange to make room for the dissolved species.
- This usually has its own enthalpy cost.
- Organizing solvent can decrease disorder.
Solute-solvent interaction (solvation)
- Water surrounds and interacts with the dissolved particles.
- This usually releases energy (exothermic enthalpy contribution).
- The effect on disorder depends on how structured the solvation shells become.
The overall free energy combines all three contributions, and each one depends on both enthalpy and entropy:
ΔG° = ΔH° - TΔS°
The important takeaway is that you can estimate the sign and relative magnitude of the enthalpy and entropy parts of each factor, but the total ΔG° is hard to predict because the factors often cancel.
The table below summarizes typical trends for each step. Treat it as a reasoning guide, not a rule that always holds.
| Process | Enthalpy Change (ΔH°) | Entropy Change (ΔS°) | Typical Effect on Dissolution |
|---|---|---|---|
| Breaking the solid apart | Positive (endothermic) | Positive (more disorder) | Entropy tends to favor dissolution |
| Reorganizing the solvent | Often positive | Negative (less disorder) | Tends to oppose dissolution |
| Solvating the species | Negative (exothermic) | Often negative | Enthalpy tends to favor dissolution |
The Role of Enthalpy
The overall enthalpy change tells you whether dissolution absorbs or releases heat. For an ionic compound, it is useful to think of two big competing pieces:
- The energy to break the ionic lattice, which costs energy.
- The energy released when ions are surrounded by water (hydration), which releases energy.
When dissolution is exothermic (ΔH° < 0):
- Heat is released and the solution warms.
- The hydration contribution outweighs the cost of breaking the lattice.
When dissolution is endothermic (ΔH° > 0):
- Heat is absorbed and the solution cools.
- The cost of breaking the lattice outweighs the hydration release.
What makes these enthalpy pieces large or small:
For breaking the solid:
- Smaller ions and higher charges form stronger lattices that are harder to break.
For hydrating the ions:
- Smaller ions and higher charges interact more strongly with water, releasing more energy.
Notice the trap built in: small, highly charged ions make both the lattice and the hydration stronger. Those two large effects push in opposite directions, which is exactly why total predictions are hard.
The Role of Entropy
Entropy change measures how disorder shifts during dissolution, and it also reflects all three steps.
- Breaking up an ordered solid increases disorder (positive contribution).
- Building organized solvent shells around ions decreases disorder (negative contribution).
For ionic compounds:
- Small, highly charged ions force a lot of water organization, giving a large negative entropy contribution.
- Larger ions with spread-out charge organize water less, so the entropy change can be smaller or even positive.
Because the entropy pieces also partly cancel, the sign of the total entropy change is not always obvious from one factor alone.
Predicting Solubility Trends
Once you have the signs of ΔH° and ΔS°, you can reason about whether dissolution is favored and how temperature changes it.
For dissolution to be thermodynamically favored:
ΔG° = ΔH° - TΔS° < 0
| Enthalpy (ΔH°) | Entropy (ΔS°) | Result | Temperature Effect |
|---|---|---|---|
| Negative | Positive | Favored at all T | Often less soluble at higher T |
| Negative | Negative | Favored only if |ΔH°| > T|ΔS°| | Often less soluble at higher T |
| Positive | Positive | Favored only if T|ΔS°| > |ΔH°| | More soluble at higher T |
| Positive | Negative | Not favored at any T | Effectively insoluble |
This is the same sign-analysis you use for any process with ΔG° = ΔH° - TΔS°, applied to dissolving a salt.
Why Predictions Are Hard
The principles look clean, but real predictions get messy because the contributions often cancel:
- Lattice and hydration enthalpies are both large but opposite in sign, so a small change in either can flip the overall enthalpy.
- Ion size and charge raise both lattice strength and hydration strength, creating competing effects.
- Solvent properties matter: polar solvents interact more strongly with ions, and the solvent's polarity shapes how well ions are stabilized.
This is the heart of the topic: you can describe each piece, but the model only goes so far in predicting the total free energy of dissolution.
How to Use This on the AP Chemistry Exam
Free Response
- Explain solubility by linking it to enthalpy and entropy changes during dissolution, not just by stating ΔG° = ΔH° - TΔS°.
- When asked about temperature effects, reason from the signs of ΔH° and ΔS°. Endothermic dissolution with a positive entropy change becomes more favored as T rises.
- If asked to evaluate a model, point out that the three contributions to free energy often cancel, so estimating each piece is easier than predicting the total.
Problem Solving
- Identify the sign of each contribution first: breaking the solid, reorganizing solvent, solvation.
- Combine the enthalpy pieces and the entropy pieces separately, then apply ΔG° = ΔH° - TΔS°.
- Connect a negative ΔG° to a favored dissolution, and remember that solubility is the macroscopic result.
Common Trap
- Do not assume an exothermic dissolution must be more favored than an endothermic one. Entropy can drive an endothermic salt to dissolve.
- Do not treat "small, highly charged ion" as automatically more soluble or less soluble. It strengthens both lattice and hydration, so the net result is not obvious.
Common Misconceptions
- "If a salt dissolves and the solution gets cold, dissolution is not favored." A cold solution just means dissolution is endothermic. A positive entropy change can still make ΔG° negative, so it is favored.
- "Stronger hydration always means more soluble." Stronger hydration helps, but the same features that boost hydration also strengthen the lattice. You have to weigh both.
- "You can plug numbers in and reliably predict ΔG° of dissolution." Because the enthalpy and entropy contributions often nearly cancel, the total is hard to predict precisely, even when you know each step's general trend.
- "Entropy always increases when a solid dissolves." Breaking the solid increases disorder, but organizing water around the ions decreases it. The overall entropy change can be small or even negative.
- "Temperature affects every salt the same way." Whether solubility rises or falls with temperature depends on the signs of ΔH° and ΔS° for that specific dissolution.
Related AP Chemistry Guides
Vocabulary
The following words are mentioned explicitly in the College Board Course and Exam Description for this topic.Term | Definition |
|---|---|
dissolution | The process by which a solute dissolves in a solvent to form a solution, involving the breaking of bonds or interactions in the solute and formation of new interactions with the solvent. |
enthalpy | The total heat content of a system; at constant pressure, the enthalpy change equals the thermal energy transferred to or from the surroundings during a chemical or physical process. |
entropy | A measure of the disorder or randomness in a system, including the dispersal of dissolved particles and reorganization of solvent molecules during dissolution. |
Gibbs free energy change | The change in Gibbs free energy (ΔG°) for a chemical or physical process, measured under standard conditions, that indicates whether a process is thermodynamically favored. |
intermolecular interactions | Forces between molecules, such as hydrogen bonding, dipole-dipole forces, and London dispersion forces, that affect the physical and chemical properties of substances. |
salt | An ionic compound formed from the reaction of an acid and a base. |
solubility | The maximum amount of a solute that can dissolve in a given amount of solvent at a specific temperature, typically expressed in moles per liter (molarity) or grams per 100 mL of solvent. |
solvent | The substance, typically a liquid, in which a solute dissolves to form a solution. |
Frequently Asked Questions
What is AP Chemistry 9.6 about?
AP Chemistry 9.6 is about free energy of dissolution and salt solubility. You explain how enthalpy and entropy changes during dissolution affect whether a salt dissolves and why predicting the overall free energy change can be difficult.
What is free energy of dissolution?
Free energy of dissolution is the overall Gibbs free energy change for a substance dissolving. It reflects particle-level changes from breaking interactions in the solid, reorganizing the solvent, and forming solute-solvent interactions.
What factors affect ΔG° for dissolution?
Three main factors affect ΔG° for dissolution: breaking the interactions that hold the solid together, reorganizing solvent around dissolved species, and interactions between dissolved species and the solvent. Each factor has enthalpy and entropy contributions.
Why is solubility hard to predict from ΔH° and ΔS°?
Solubility can be hard to predict because the enthalpic and entropic contributions from the three dissolution steps often partly cancel. A simple model may estimate signs or relative sizes, but the total ΔG° can still be challenging to predict.
How does temperature affect dissolution?
Temperature affects dissolution through $\Delta G^\circ = \Delta H^\circ - T\Delta S^\circ$. Whether higher temperature makes dissolution more favorable depends on the signs and relative sizes of $\Delta H^\circ$ and $\Delta S^\circ$.
How should I answer AP Chem free energy of dissolution questions?
Connect the macroscopic observation, such as solubility, to particle-level interactions and thermodynamic reasoning. Name the relevant enthalpy and entropy changes, then explain why cancellations among factors can limit simple predictions.