Thermodynamics

The One Thing

Thermodynamics describes how energy is transferred and transformed during physical and chemical processes while conserving total energy. Enthalpy, entropy, and Gibbs energy determine heat changes, the natural direction of processes, and their thermodynamic feasibility.

Reactions, Processes and Experiments

Key Terms

• System: The part of the universe selected for study.
• Surroundings: Everything outside the system that can exchange energy or matter with it.
• Boundary: The real or imaginary surface separating the system from its surroundings.
• 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 that depends only on the initial and final states, not on the path followed.
• Path function: A quantity whose value depends on the route taken during a process, such as heat or work.
• Internal energy: The total microscopic energy of a system, including molecular kinetic and potential energies.
• Heat: Energy transferred between a system and surroundings because of a temperature difference.
• Work: Energy transferred when a force causes displacement; in chemistry, expansion or compression work is important.
• First law of thermodynamics: Energy cannot be created or destroyed; it can only be transferred or transformed.
• Enthalpy: A state function defined as $H=U+PV$; its change represents heat absorbed or released at constant pressure.
• Exothermic process: A process that releases heat to the surroundings and has $\Delta H<0$.
• Endothermic process: A process that absorbs heat from the surroundings and has $\Delta H>0$.
• Calorimetry: The measurement of heat changes during physical or chemical processes.
• Specific heat capacity: The heat required to raise the temperature of one unit mass of a substance by 1 K.
• Molar heat capacity: The heat required to raise the temperature of one mole of a substance by 1 K.
• Entropy: A measure of the randomness, disorder, or dispersal of energy in a system.
• Second law of thermodynamics: The entropy of an isolated system tends to increase for a spontaneous process.
• Spontaneous process: A process that can occur on its own under given conditions, without continuous external assistance.
• Gibbs energy: A thermodynamic quantity defined as $G=H-TS$; its change predicts spontaneity at constant temperature and pressure.
• Standard enthalpy of formation: The enthalpy change when one mole of a compound forms from its elements in their standard states.
• Hess's law: The total enthalpy change is the same whether a reaction occurs in one step or several steps.

Easily Confused

• Heat and work: Both are path functions representing energy transfer; heat results from a temperature difference, whereas work involves energy transfer by force and displacement.
• Internal energy and enthalpy: Both are state functions; internal energy is $U$, whereas enthalpy is defined by $H=U+PV$.
• Exothermic and endothermic processes: Exothermic processes release heat and have $\Delta H<0$; endothermic processes absorb heat and have $\Delta H>0$.
• Spontaneity and reaction rate: A spontaneous process is thermodynamically feasible, but it may be very slow because of a high activation energy.
• Constant-volume and constant-pressure conditions: At constant volume, $q_v=\Delta U$ because $w=0$; at constant pressure, $q_p=\Delta H$.
• System entropy and universe entropy: A spontaneous process requires $\Delta S_{\text{universe}}>0$, not necessarily $\Delta S_{\text{system}}>0$.
• Calorimeter heat and reaction heat: $q_{\text{reaction}}=-q_{\text{surroundings}}$, so the reaction and surroundings have opposite heat changes.
• Gibbs energy and enthalpy: Enthalpy indicates heat release or absorption, whereas Gibbs energy combines enthalpy and entropy to assess spontaneity at constant temperature and pressure.

What Gets Asked

• Define and distinguish open, closed, and isolated systems; marks depend on identifying whether matter, energy, or neither is exchanged.
• Apply the first-law equation $\Delta U=q+w$, including the chemistry sign convention: $q>0$ for absorbed heat and $q<0$ for released heat.
• Calculate expansion or compression work using $w=-P_{\text{ext}}\Delta V$, identifying expansion as $w<0$ and compression as $w>0$.
• Use calorimetry equations, including $q=mc\Delta T$, $q_{\text{calorimeter}}=C_{\text{calorimeter}}\Delta T$, and $q_{\text{reaction}}=-q_{\text{surroundings}}$; the common error is assigning the same sign to reaction and surroundings.
• Calculate reaction enthalpy using Hess's law or ${\Delta }_{r}H°=\Sigma {\Delta }_{f}H°(\text{products})-\Sigma {\Delta }_{f}H°(\text{reactants})$, remembering that reversing an equation changes the sign of $\Delta H$ and multiplying an equation multiplies $\Delta H$ by the same factor.
• Determine spontaneity using $\Delta G=\Delta H-T\Delta S$, with $T$ in kelvin, and distinguish $\Delta G<0$, $\Delta G=0$, and $\Delta G>0$ from reaction speed.