Unit 6: Thermodynamics

Thermodynamics

The study of how energy is transferred and transformed; used in chemistry to predict heat flow and spontaneity of processes.

System

The part of the universe chosen for study (often the reacting chemicals).

Surroundings

Everything outside the system that can exchange energy (and sometimes matter) with it.

Open system

A system that can exchange both matter and energy with the surroundings.

Closed system

A system that can exchange energy but not matter with the surroundings.

Isolated system

A system that exchanges neither matter nor energy with the surroundings (idealized).

Temperature

A measure of the average kinetic energy of particles (molecular motion).

Heat

Energy transferred between substances due to a temperature difference (energy in transit).

Endothermic process

A process in which the system absorbs heat from the surroundings; typically ΔH > 0.

Exothermic process

A process in which the system releases heat to the surroundings; typically ΔH < 0.

Enthalpy (H)

A thermodynamic quantity defined as H = E + PV; often used because many reactions occur at constant pressure.

Enthalpy change (ΔH)

The change in enthalpy between products and reactants; negative for exothermic reactions and positive for endothermic reactions.

First Law of Thermodynamics

Energy is conserved: it cannot be created or destroyed, only transferred or transformed.

Internal energy change (ΔE)

Change in a system’s internal energy; ΔE = q + w.

Heat sign convention (q)

q is positive when heat flows into the system and negative when heat leaves the system.

Work sign convention (w)

w is positive when work is done on the system and negative when the system does work on the surroundings.

Pressure–volume work

Work associated with expansion/compression: w = −P_extΔV (expansion gives w < 0; compression gives w > 0).

Reaction coordinate (energy) diagram

A plot of potential energy versus reaction progress used to visualize activation energy and ΔH.

Activation energy (E_a)

The energy barrier from reactants up to the transition state on an energy diagram.

Transition state

The highest-energy point along the reaction pathway (top of the energy barrier).

Catalyst

A substance that speeds up a reaction by lowering E_a; it does not change reactant/product energies, ΔH, ΔG, ΔS, K, or the equilibrium position.

Calorimetry

The measurement of heat transfer by observing temperature changes.

Heat capacity

The amount of heat required to raise an object’s temperature by 1 K (or 1 °C).

Specific heat capacity (c)

Heat required to raise 1 gram of a substance by 1 K (or 1 °C); units often J g⁻¹ °C⁻¹.

Calorimetry equation (q = mcΔT)

Calculates heat gained/lost by a material: q = mc(Tf − Ti).

Coffee-cup calorimeter

A constant-pressure calorimeter (often styrofoam cup) used to measure heat of reaction in solution.

Constant-pressure heat (q_p = ΔH)

At constant pressure, the heat transferred equals the enthalpy change: q_p = ΔH.

Energy conservation in calorimetry

For system + surroundings: qsys + qsurr = 0.

Reaction heat sign in coffee-cup calorimetry

Heat of reaction is the opposite sign of the surroundings: qrxn = −qsurr (often q_surr = mcΔT for the solution).

Bomb calorimeter

A rigid, sealed constant-volume calorimeter; the heat measured corresponds to ΔE (not ΔH).

Constant-volume heat (q_v = ΔE)

At constant volume, heat transfer equals the change in internal energy: q_v = ΔE.

Calorimeter heat capacity (C_cal)

Overall heat capacity of a calorimeter; qcal = CcalΔT.

Relation between ΔH and ΔE (gases)

For reactions involving gases: ΔH = ΔE + Δn_gasRT.

State function

A property that depends only on initial and final states, not the path taken (e.g., H, S, G).

Hess’s Law

Because ΔH is a state function, enthalpy changes add when reactions are added to obtain a target reaction.

Thermochemical equation rules

If a reaction is reversed, ΔH changes sign; if coefficients are multiplied, ΔH is multiplied by the same factor (states matter).

Heating curve

Graph showing temperature vs heat added; temperature rises within a phase, but stays constant during phase changes while energy changes potential energy.

Phase change enthalpies

ΔHfus is energy to melt at the melting point; ΔHvap is energy to vaporize at the boiling point; the reverse processes release the same magnitude of energy.

Enthalpy of solution (ΔH_soln)

Overall enthalpy change when an ionic solid dissolves; includes energy absorbed to separate ions and solvent molecules and energy released from ion–dipole (hydration) interactions (hydration is always negative and stronger for higher charge/smaller ions).

Standard state

Reference condition for thermodynamic data: pure solids/liquids in most stable form at 1 bar; gases at 1 bar; solutes at 1.0 M (commonly at 298 K).

Standard enthalpy of formation (ΔH_f°)

Enthalpy change to form 1 mole of a compound from its elements in their standard states; elements in standard states have ΔH_f° = 0.

Formation reaction (standard)

A balanced equation that forms exactly 1 mole of product from elements in their standard states; may use fractional coefficients (e.g., 1/2 O₂).

Reaction enthalpy from formation enthalpies

ΔHrxn° = Σ nΔHf°(products) − Σ nΔH_f°(reactants).

Bond enthalpy (bond energy)

Energy required to break 1 mole of a specific bond in gaseous molecules; typically given as average values, so calculations are approximate.

Estimating ΔH using bond enthalpies

ΔH_rxn ≈ ΣD(bonds broken) − ΣD(bonds formed); break = energy in, form = energy out.

Entropy (S)

A measure of energy dispersal and the number of accessible microstates; generally: gas > liquid > solid, and mixing/increasing temperature tends to increase S.

Standard molar entropy and reaction entropy

Standard molar entropies (S°) are not zero for elements; compute ΔS_rxn° = Σ nS°(products) − Σ nS°(reactants).

Second Law of Thermodynamics

Spontaneous processes increase total entropy: ΔSuniv = ΔSsys + ΔSsurr, and spontaneity requires ΔSuniv > 0 (at equilibrium, ΔS_univ = 0).

Gibbs free energy (ΔG)

Criterion for spontaneity at constant T and P: ΔG = ΔH − TΔS; ΔG < 0 spontaneous, ΔG = 0 equilibrium, ΔG > 0 nonspontaneous.

Free energy & equilibrium (K and Q)

ΔG° = −RT ln K; ΔG = ΔG° + RT ln Q. If Q < K the reaction proceeds forward, if Q > K it proceeds backward, and at equilibrium Q = K with ΔG = 0.