AP Chemistry Thermodynamics- Essential Concepts and Study Guide
What You Need to Know About Thermodynamics on the AP Chemistry Exam
Thermodynamics shows up in roughly 25-30% of AP Chemistry free-response questions. You cannot afford to bomb this unit. The good news? The concepts are straightforward once you stop overcomplicating things.
The Big Three: Internal Energy, Enthalpy, and Entropy
These three concepts form the foundation. Everything else in thermodynamics builds on them.
Internal Energy (U)
Internal energy is the total energy contained within a system. It includes kinetic energy of molecules and potential energy from chemical bonds. You cannot measure absolute internal energy directly. You measure changes in internal energy.
The equation is:
ΔU = q + w
Where:
- ΔU = change in internal energy
- q = heat transferred
- w = work done on or by the system
Pay attention to sign conventions. Heat entering the system is positive. Work done on the system is positive. Work done by the system is negative (like gas expansion).
Enthalpy (H)
Enthalpy is heat content at constant pressure. The AP exam almost always uses enthalpy because most chemistry happens in open containers (constant atmospheric pressure).
ΔH = ΔU + PΔV
For most reactions, you just need to know that ΔH > 0 means endothermic (heat absorbed) and ΔH < 0 means exothermic (heat released).
Entropy (S)
Entropy measures disorder or randomness in a system. The universe tends toward disorder, which is why entropy is a key concept in predicting spontaneity.
Key points about entropy:
- Gases have higher entropy than liquids
- Liquids have higher entropy than solids
- More moles of gas = higher entropy
- Temperature increases generally increase entropy
The Laws of Thermodynamics
First Law: Energy Conservation
Energy cannot be created or destroyed, only transferred. This is why the universe's total energy stays constant. Your calculations must balance.
Second Law: Entropy and Spontaneity
For a process to be spontaneous, the total entropy of the universe must increase. This does not mean the entropy of your system must increase—only the universe's overall entropy.
A reaction can have decreasing entropy (becoming more ordered) if the surroundings release enough heat to compensate.
Third Law: Absolute Entropy
A perfect crystal at absolute zero has zero entropy. This gives us reference points for calculating absolute entropies of substances. You will rarely need this directly on the exam, but it explains why standard molar entropy (S°) values exist.
Gibbs Free Energy: The Real Predictor of Spontaneity
Gibbs free energy combines enthalpy, entropy, and temperature into one equation that tells you whether a reaction is spontaneous.
ΔG = ΔH - TΔS
| ΔH | ΔS | ΔG | Reaction Type |
|---|---|---|---|
| Negative | Positive | Always negative | Always spontaneous |
| Positive | Negative | Always positive | Never spontaneous |
| Negative | Negative | Negative at low T | Spontaneous at low temperature |
| Positive | Positive | Negative at high T | Spontaneous at high temperature |
This table is the backbone of Gibbs free energy questions. Memorize it.
Hess's Law: Calculating Enthalpy Changes
Hess's Law states that enthalpy change is independent of the reaction path. You can add reactions together to get the overall enthalpy change.
Rules for Hess's Law problems:
- Reverse a reaction? Flip the sign of ΔH
- Multiply a reaction? Multiply ΔH by that coefficient
- Add reactions together? Add their ΔH values
Example: If you need to find ΔH for 2A → C, and you know:
- A → B, ΔH = -50 kJ
- B → C, ΔH = -30 kJ
Multiply the first reaction by 2, then add them. Your answer is (-50 × 2) + (-30) = -130 kJ.
Calorimetry: Measuring Heat
Calorimetry experiments measure heat flow using the relationship:
q = mcΔT
Where:
- q = heat (J or kJ)
- m = mass (g)
- c = specific heat capacity (J/g·°C)
- ΔT = change in temperature (final - initial)
Water has a specific heat of 4.184 J/g·°C. This is one of the most commonly used values on the exam.
In coffee cup calorimetry (constant pressure), the heat absorbed by the solution equals the heat released by the reaction (with opposite sign). In bomb calorimetry (constant volume), you account for the calorimeter's heat capacity.
Standard Enthalpy of Formation (ΔH°f)
The standard enthalpy of formation is the enthalpy change when one mole of a compound forms from its elements in their standard states.
Key facts:
- ΔH°f for any element in its standard state = 0
- Use a formation equation: elements → compound
- Tabulated values are per mole of compound formed
To calculate ΔH° for a reaction using formation data:
ΔH° = Σ(n × ΔH°f products) - Σ(n × ΔH°f reactants)
Bond Enthalpies: An Alternative Method
You can estimate reaction enthalpy by breaking all bonds in reactants (endothermic) and forming all bonds in products (exothermic).
ΔH ≈ Σ(bonds broken) - Σ(bonds formed)
Bond enthalpies give approximate values. Formation data gives exact values. The AP exam may ask you to use either method, so know both.
Getting Started: Your Study Approach
Follow these steps to master AP Chemistry Thermodynamics:
- Memorize the core equations: ΔU = q + w, ΔG = ΔH - TΔS, q = mcΔT, and Hess's Law operations
- Practice sign conventions until they become automatic—half of all thermodynamics errors come from sign mistakes
- Work through past FRQs from College Board—thermodynamics appears every year
- Memorize the Gibbs free energy table above—it appears in some form on almost every exam
- Understand why, not just what: knowing that entropy increases with temperature matters more than memorizing definitions
Common Mistakes That Cost Points
- Forgetting to flip the sign when reversing a reaction in Hess's Law
- Confusing enthalpy units (kJ vs kJ/mol)
- Not converting temperatures to Kelvin in Gibbs free energy calculations
- Forgetting that gases have much higher entropy than solids or liquids
- Mixing up endothermic/exothermic sign conventions
What Actually Shows Up on the Exam
Based on recent AP Chemistry exams, expect:
- Calculating ΔH from calorimetry data
- Using Hess's Law to find unknown enthalpy changes
- Determining spontaneity with Gibbs free energy
- Connecting thermodynamics to equilibrium (ΔG = -RT ln K)
- Interpreting heating/cooling curves
The connection between thermodynamics and equilibrium常数 is particularly important. Remember: ΔG° = -RT ln K. When ΔG° is negative, K > 1. When ΔG° is positive, K < 1.