Chemical Thermodynamics

The One Thing

Chemical thermodynamics relates energy conservation to heat, work, enthalpy, entropy, and spontaneity. At constant temperature and pressure, Gibbs free energy provides the principal criterion: $\Delta G<0$ indicates a spontaneous forward process, $\Delta G>0$ a non-spontaneous forward process, and $\Delta G=0$ equilibrium.

Reactions, Processes and Experiments

Key Terms

• System: The part of the universe selected for study, such as a reacting mixture.
• Surroundings: Everything outside the system that can exchange energy or matter with it.
• Open system: A system that exchanges both matter and energy with its surroundings.
• Closed system: A system that exchanges energy but not matter with its surroundings.
• Isolated system: A system that exchanges neither matter nor energy with its surroundings.
• State function: A property whose change depends only on the initial and final states, not on the path followed.
• Extensive property: A property that depends on the amount of substance, such as mass, volume, and internal energy.
• Intensive property: A property independent of the amount of substance, such as temperature, pressure, and density.
• Internal energy ($U$): The total microscopic kinetic and potential energy of particles in a system.
• Heat ($q$): Energy transferred between a system and its surroundings because of a temperature difference.
• Work ($w$): Energy transferred when an external force acts on or through the system, such as expansion work.
• First law of thermodynamics: Energy cannot be created or destroyed; it can only be transferred or converted from one form to another.
• Enthalpy ($H$): A state function defined as $H=U+pV$; its change represents heat absorbed or released at constant pressure when only pressure-volume work occurs.
• Exothermic process: A process that releases heat to the surroundings, giving a negative enthalpy change.
• Endothermic process: A process that absorbs heat from the surroundings, giving a positive enthalpy change.
• Entropy ($S$): A state function related to the dispersal of energy and the number of possible microscopic arrangements.
• Second law of thermodynamics: A spontaneous process increases the total entropy of the universe.
• Gibbs free energy ($G$): A thermodynamic function defined as $G=H-TS$; its change predicts spontaneity at constant temperature and pressure.
• Spontaneous process: A process that can occur on its own under specified conditions, although it may be slow.
• Equilibrium: The state in which the forward and reverse processes occur at equal rates and the macroscopic properties remain constant.
• Hess's law: The total enthalpy change of a reaction is the same whether the reaction occurs in one step or several steps.
• Standard enthalpy change: The enthalpy change measured under standard conditions, commonly with substances in their standard states and a pressure of 1 bar.

Easily Confused

• Heat and work: Heat is transferred because of a temperature difference; work is transferred when an external force acts on or through the system.
• Internal energy and enthalpy: Internal energy is $U$, whereas enthalpy is $H=U+pV$ and is especially useful for constant-pressure reactions.
• Exothermic and endothermic processes: Exothermic processes have negative $\Delta H$ and release heat; endothermic processes have positive $\Delta H$ and absorb heat.
• Spontaneity and reaction rate: $\Delta G<0$ concerns thermodynamic feasibility, whereas reaction rate depends mainly on kinetics and activation energy.
• State functions and path-dependent transfers: State-function changes depend only on initial and final states; heat and work describe transfers along a process path.
• Open, closed, and isolated systems: Open systems exchange matter and energy; closed systems exchange energy but not matter; isolated systems exchange neither.
• Extensive and intensive properties: Extensive properties depend on the amount of substance; intensive properties do not.
• $\Delta U$ and $\Delta H$: At constant volume, $q_v=\Delta U$; at constant pressure with only pressure-volume work, $q_p=\Delta H$.
• Entropy of the system and entropy of the universe: A spontaneous process requires $\Delta S_{\text{universe}}>0$, not necessarily $\Delta S_{\text{system}}>0$.
• Hess's law and bond-enthalpy estimates: Hess's law uses exact addition of known thermochemical equations, whereas bond enthalpies provide an estimate.

What Gets Asked

• Sign-convention questions: Apply $\Delta U=q+w$, taking heat absorbed by the system and work done on the system as positive. Reversing the sign of expansion work loses marks.
• Constant-volume and constant-pressure calculations: Identify whether $q_v=\Delta U$ or $q_p=\Delta H$ applies before substituting values.
• Calorimetry calculations: Use $q=mc\Delta T$ with mass, specific heat capacity, and temperature change; confusing $c$ with $C$ or $C_m$ gives an incorrect result.
• Hess's law and formation-enthalpy calculations: Add or reverse thermochemical equations correctly and use ${\Delta }_r{H}^{\circ }=\Sigma \nu {\Delta }_f{H}^{\circ }(\text{products})-\Sigma \nu {\Delta }_f{H}^{\circ }(\text{reactants})$. The standard enthalpy of formation of an element in its most stable standard state is zero.
• Entropy and spontaneity questions: Apply $\Delta G=\Delta H-T\Delta S$, use kelvin for temperature, and distinguish $\Delta G<0$, $\Delta G>0$, and $\Delta G=0$.
• Conceptual comparisons: Explain why a process may be spontaneous even when endothermic, and why thermodynamics predicts feasibility but not how quickly a reaction occurs.