Rucete ✏ AP Chemistry In a Nutshell
11. Thermodynamics
This chapter explores energy changes in chemical systems, focusing on heat, work, entropy, and free energy. It introduces core thermodynamic concepts and equations that govern spontaneity and equilibrium.
Systems and State Functions
• A system is the part of the universe under study; surroundings are everything else.
• Open systems exchange energy and matter; closed systems exchange only energy; isolated systems exchange neither.
• State functions depend only on the initial and final states, not the path (e.g., pressure, volume, temperature, ΔE, ΔH, ΔS, ΔG).
• Heat (q) and work (w) are not state functions since they depend on the process path.
Standard State and Sign Conventions
• Standard state: 1 atm pressure, 25°C (298 K), 1 mole of substance.
• Thermodynamic quantities under standard conditions use superscript “°” (e.g., ΔH°, ΔG°).
• Exothermic processes release energy (ΔH < 0); endothermic processes absorb energy (ΔH > 0).
Forms of Energy
• Kinetic energy (KE): energy of motion, KE = ½mv².
• Potential energy (PE): stored energy, such as electrostatic or gravitational energy.
• Total energy (E) = PE + KE.
• Energy is conserved and can be converted between forms.
Specific Heat and Calorimetry
• Specific heat: energy required to raise the temperature of 1 g of a substance by 1°C.
• For water, specific heat = 4.184 J/g°C.
• Heat energy is calculated using q = mcΔT.
• In calorimetry, heat lost by one substance is gained by another: q₁ + q₂ = 0.
First Law of Thermodynamics
• Energy cannot be created or destroyed; ΔE = q + w.
• Positive q: heat added to system. Positive w: work done on system.
• Work (w) = –PΔV; expansion does work on surroundings (w < 0), compression does work on system (w > 0).
Enthalpy (ΔH)
• ΔH = qp at constant pressure; it represents the heat content of a reaction.
• ΔH = Hproducts – Hreactants
• At constant volume (in a bomb calorimeter), ΔE = qv, since PΔV = 0.
• For most reactions, ΔH ≈ ΔE.
Hess’s Law
• If equations are added, their ΔH values are also added.
• Multiplying an equation by a factor multiplies ΔH by the same factor.
• Reversing an equation changes the sign of ΔH.
Standard Enthalpies of Formation (ΔH°f)
• ΔH°f: enthalpy change when 1 mol of a compound forms from its elements in their standard states.
• ΔH°rxn = ΣnΔH°f(products) – ΣnΔH°f(reactants)
• Elements in their standard states (e.g., O₂, H₂, N₂) have ΔH°f = 0.
Bond Enthalpies
• Bond enthalpy: energy required to break 1 mol of a bond in the gas phase.
• ΔH°rxn ≈ Σ(bonds broken) – Σ(bonds formed)
• Breaking bonds requires energy (positive ΔH); forming bonds releases energy (negative ΔH).
Entropy (S)
• Entropy is a measure of disorder or randomness.
• ΔS = Sfinal – Sinitial
• Processes that increase disorder (e.g., melting, vaporization, mixing) increase entropy.
• The second law of thermodynamics: total entropy of the universe increases in a spontaneous process (ΔSuniv > 0).
Factors Affecting Entropy
• More particles → higher entropy.
• Gas > liquid > solid in entropy.
• Higher temperature → higher entropy.
• Dissolving solids or liquids generally increases entropy.
Third Law of Thermodynamics
• Entropy of a perfect crystalline substance at 0 K is zero.
• Allows for absolute entropy (S°) values in J/mol·K for all substances.
Gibbs Free Energy (ΔG)
• ΔG = ΔH – TΔS (at constant temperature).
• Determines spontaneity: – ΔG < 0 → spontaneous – ΔG > 0 → nonspontaneous – ΔG = 0 → at equilibrium
• Combines enthalpy and entropy into a single value.
Free Energy and Equilibrium
• At equilibrium: ΔG = 0 and ΔG° = –RT ln K
• Large K → negative ΔG°, reaction favors products.
• Small K → positive ΔG°, reaction favors reactants.
Temperature and Spontaneity
• If ΔH < 0 and ΔS > 0 → always spontaneous.
• If ΔH > 0 and ΔS < 0 → never spontaneous.
• If both ΔH and ΔS are positive → spontaneous at high T.
• If both ΔH and ΔS are negative → spontaneous at low T.
In a Nutshell
Thermodynamics studies energy changes and predicts reaction spontaneity using enthalpy, entropy, and free energy. The laws of thermodynamics guide how heat and work flow, while Gibbs free energy indicates whether a reaction will occur naturally and how it relates to equilibrium and temperature conditions.