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Credit Sprint's Guide to Chemistry Unit 6: Thermodynamics

AP® weight 7–9%CLEP® weight 4–6%

Thermodynamics tracks energy changes. Heat is energy moved by a temperature difference, enthalpy is the heat at constant pressure, and Hess's law, formation enthalpies and bond energies let you calculate ΔH without measuring it.

Heating curveHeat addedTemperature melting (ΔHfus)boiling (ΔHvap) solidliquidgas

What you need to know

  • Temperature measures average kinetic energy; heat q is energy transferred because of a temperature difference.
  • q = mcΔT. Heat lost by the hot object equals heat gained by the cold one at equilibrium.
  • Exothermic: ΔH < 0, heat leaves the system, surroundings warm up. Endothermic: ΔH > 0.
  • During a phase change temperature is constant; energy changes the potential energy of the particles (q = nΔHfus or nΔHvap).
  • Hess's law: ΔH is a state function. Reverse an equation and the sign flips; multiply an equation and ΔH multiplies.
  • ΔH°rxn = ΣΔH°f(products) − ΣΔH°f(reactants). ΔH°f of an element in its standard state is zero.
  • Bond energies: ΔH ≈ Σ(bonds broken) − Σ(bonds formed). Breaking bonds is endothermic.
  • Calorimetry measures q for a reaction from the temperature change of a known mass of water.
  • First law: ΔU = q + w; energy is conserved.

Equations and rules

q = mcΔT
q = nΔH(phase change)
ΔH°rxn = ΣΔH°f(products) − ΣΔH°f(reactants)
ΔH ≈ Σ(bonds broken) − Σ(bonds formed)

Exam traps

Watch forReversing a reaction changes the sign of ΔH.
Watch forThe coffee-cup calorimeter measures the surroundings, so q_rxn = −q_solution.
Watch forBond energy calculations are only estimates.
Watch forTemperature doesn't change during melting or boiling.

Key terms

  • Enthalpy change (ΔH): Heat transferred at constant pressure; negative for exothermic reactions.
  • Specific heat capacity: Heat needed to raise 1 g of a substance by 1 °C; q = mcΔT.
  • Hess's law: The enthalpy change of a reaction equals the sum of the enthalpy changes of its steps.
  • Standard enthalpy of formation: Enthalpy change to form 1 mol of a compound from its elements in standard states.
  • Bond energy: Energy needed to break one mole of a bond; breaking is endothermic.
  • Calorimetry: Measuring heat transfer from the temperature change of a known mass.
  • State function: A property that depends only on the current state, like enthalpy or entropy.
  • First law of thermodynamics: Energy is conserved: ΔU = q + w.

Practice questions

Five of the 44 Unit 6 questions in Credit Sprint. Every choice is explained in the app.

1. How much heat is needed to raise the temperature of 50.0 g of water from 25.0 °C to 35.0 °C? (c = 4.184 J/g·°C)
  1. 2090 J
  2. 5230 J
  3. 20,900 J
  4. 209 J
Show answer

2090 J. q = mcΔT = (50.0)(4.184)(10.0) = 2092 J.

2. Which expression correctly states Hess's law?
  1. ΔH of a reverse reaction has the same sign
  2. ΔH depends on the path taken
  3. The enthalpy change of a reaction is the sum of the enthalpy changes of its steps
  4. ΔH is always negative
Show answer

The enthalpy change of a reaction is the sum of the enthalpy changes of its steps. ΔH is a state function.

3. Breaking a chemical bond is
  1. endothermic
  2. neither
  3. spontaneous
  4. exothermic
Show answer

endothermic. Energy must be supplied.

4. A system absorbs 500 J of heat and does 200 J of work on the surroundings. What is the change in internal energy?
  1. +300 J
  2. +700 J
  3. −300 J
  4. +200 J
Show answer

+300 J. ΔU = q + w = 500 − 200 = +300 J.

5. If a reaction has ΔH = −100 kJ, what is ΔH for the reverse reaction?
  1. +100 kJ
  2. −100 kJ
  3. −50 kJ
  4. 0 kJ
Show answer

+100 kJ. Reversing a reaction changes the sign.