📚Study Guide: Thermodynamics
Unit 6: Thermodynamics
Thermodynamics governs the energy changes that accompany chemical and physical processes, providing the ultimate criterion for whether a reaction will occur spontaneously. This unit focuses on enthalpy (delta H), entropy (delta S), and Gibbs free energy (delta G), and the relationships among them. Enthalpy changes are measured experimentally using calorimetry or calculated using Hess's law and standard enthalpies of formation. Entropy, a measure of molecular disorder or dispersal of energy, increases for processes that spread out matter and energy. The second law of thermodynamics states that the total entropy of the universe increases for any spontaneous process. However, chemists most often use Gibbs free energy--delta G = delta H - T delta S--to predict spontaneity at constant temperature and pressure: a negative delta G indicates a spontaneous process. Students must be able to predict the sign of delta S for various processes and understand how temperature affects the spontaneity of reactions with different signs of delta H and delta S. The AP exam frequently tests calorimetry calculations, Hess's law problems, and the interpretation of thermodynamic data tables. Understanding bond enthalpies allows estimation of reaction enthalpies by comparing the energy required to break bonds in reactants with the energy released when bonds form in products. Mastery of these concepts is essential for predicting reaction feasibility and understanding bioenergetics.
Key Concepts
- Enthalpy (delta H): The heat absorbed or released at constant pressure. Endothermic reactions absorb heat (delta H > 0); exothermic reactions release heat (delta H < 0). Calorimetry measures q = m c delta T; at constant pressure, q = delta H.
- Hess's Law: The overall enthalpy change for a reaction equals the sum of enthalpy changes for individual steps, regardless of the pathway. Allows calculation of delta H for reactions that are difficult to measure directly.
- Standard Enthalpies of Formation (delta Hf degrees): The enthalpy change when one mole of a compound forms from its elements in their standard states. delta H degrees_rxn = Sum delta Hf degrees(products) - Sum delta Hf degrees(reactants).
- Entropy (delta S): A measure of the dispersal of matter and energy. delta S increases when solids become liquids or gases, when a substance dissolves, when the number of moles of gas increases, and when temperature increases. The third law states that the entropy of a perfect crystal at absolute zero is zero.
- Gibbs Free Energy (delta G): delta G = delta H - T delta S. If delta G < 0, the process is spontaneous. If delta G > 0, it is nonspontaneous. At equilibrium, delta G = 0. Temperature determines the dominant term when delta H and delta S have the same sign.
- Bond Enthalpies: Average energy required to break one mole of a particular bond in the gas phase. delta H approximately equal to Sum(bonds broken) - Sum(bonds formed).
Vocabulary
- Calorimetry: The experimental measurement of heat flow in or out of a system during a chemical or physical process.
- State Function: A property whose value depends only on the current state of the system, not on the path taken to reach that state (e.g., enthalpy, entropy, Gibbs free energy).
- Standard State: The most stable form of a substance at 1 bar pressure and a specified temperature (usually 25 degrees C or 298 K).
- Surroundings: Everything outside the system being studied; together with the system, it constitutes the universe.
- Spontaneous Process: A process that occurs without continuous external intervention under given conditions; does NOT imply fast.
- Endothermic: A process that absorbs heat from the surroundings (delta H > 0).
Essential Formulas
- q = m c delta T
- delta H degrees_rxn = Sum delta Hf degrees(products) - Sum delta Hf degrees(reactants)
- delta S degrees_rxn = Sum S degrees(products) - Sum S degrees(reactants)
- delta G = delta H - T delta S
- delta G degrees = -R T ln K
- delta G degrees = -n F E degrees_cell
- delta H approximately equal to Sum(bonds broken) - Sum(bonds formed)
- R = 8.314 J/(mol K); F = 96,485 C/mol e-
Common Mistakes
- Confusing System and Surroundings: In calorimetry, q_system = -q_surroundings. If the solution warms up, the reaction is exothermic (releases heat).
- Assuming Exothermic Always Means Spontaneous: A reaction with delta H < 0 can be nonspontaneous at high T if delta S < 0 (delta G = delta H - T delta S becomes positive).
- Ignoring Units: When calculating delta G, convert delta H to J/mol or delta S to kJ/(mol K) so units match.
- Using Bond Enthalpies for Solids/Liquids: Bond enthalpies are defined for gases. Calculations for solids/liquids using bond enthalpies are approximations.
AP Exam Strategies
- Draw Energy Diagrams: For calorimetry and enthalpy problems, sketch the system and surroundings and indicate the direction of heat flow with arrows.
- Apply Hess's Law Systematically: Manipulate given reactions (reverse, multiply) to match the target reaction, and apply the same operations to their delta H values.
- Predict delta S Signs: When asked about entropy change, count moles of gas (increase -> delta S > 0) and consider phase changes (solid -> liquid -> gas -> delta S > 0).
- Calculate delta G at Different Temperatures: If delta H and delta S are given, calculate delta G at multiple temperatures to find where spontaneity switches sign.
Real-World Applications
- Hand Warmers: Exothermic crystallization of supersaturated sodium acetate releases heat used in reusable hand warmers.
- Metabolism: Biological systems use coupled reactions where a spontaneous process (ATP hydrolysis, delta G < 0) drives nonspontaneous biosynthesis.
- Fuel Cells: The spontaneity of hydrogen-oxygen reactions (delta G < 0) is harnessed in fuel cells to produce electricity directly.