GENERAL CHEMISTRY · THERMODYNAMICS AND THERMOCHEMISTRY STUDY MAP

Track the boundary.
Preserve the sign.

Fifteen focused outcomes connect energy ledgers, state functions, Hess cycles, entropy, Gibbs energy, calorimetry, and phase-change paths across five complete lessons and twenty original questions.

Start the lessons →Build a balanced set ↗Explore all objectives ↓
5official topics
15bounded objectives
5guided lessons
20original questions

Scope boundary

Favorability, position, and speed answer different questions.

The ADA names Laws of thermodynamics, Hess's law, Spontaneity, Enthalpies and entropies, and Heat transfer. It does not publish topic-level weights or question quotas, so these DAT TRAIN outcomes organize study without pretending to predict a test form.

Thermodynamics owns state functions, energy and entropy ledgers, spontaneity, and heat transfer. Chemical Equilibria owns Q/K position and composition; Chemical Kinetics owns rate laws, activation energy, and reaction speed. A spontaneous process can still be slow.

Energy decision sequence

Four checks before trusting an energy claim.

  1. 01

    Boundary

    Name the system, surroundings, transfer direction, and the chemical process whose energy change is being reported.

  2. 02

    Sign

    Attach every sign to the defined system and reaction direction before combining heat, work, enthalpy, or entropy terms.

  3. 03

    State or path

    Ask whether the quantity depends only on endpoints or also on how the process occurs, then choose the matching equation.

  4. 04

    Condition and units

    Check phase, pressure constraint, absolute temperature, standard state, and compatible J or kJ units before interpreting the result.

Official hierarchy → learning sequence

Five branches. Fifteen outcomes.

Open a topic to inspect its observable outcomes, must-know relationships, model boundaries, prerequisite graph, misconception corrections, representations, and direct free references.

ALaws of thermodynamics3 objectives
Objective 1GC-THM-LAW-01

System boundaries and the first law

Define a chemical system and its surroundings, then apply the chemistry sign convention in ΔE = q + w to classify heat and work crossing the boundary.

Must know
In the chemistry convention, q and w are positive when energy enters the system; expansion work done by the system is negative, while compression work done on it is positive. The first law is an energy ledger: the system's internal-energy change equals the heat and work transferred across the chosen boundary.
Model boundary
Use the sign convention and work model stated in the problem. Detailed non-PV work, irreversible-path derivations, and molecular partition functions are outside this objective unless supplied.
Misconception
Heat released by a reaction makes q positive for the reacting system. Released heat leaves the defined system, so q is negative for that system even though the surroundings gain energy.
Earlier objectives in this map
None
Objective 2GC-THM-LAW-02

State functions and path functions

Distinguish state functions from path functions and determine which energy changes depend only on initial and final states rather than on the process used.

Must know
Internal energy, enthalpy, entropy, and Gibbs free energy are state functions, so their changes depend on endpoints; heat and work describe a path between those endpoints. A thermodynamic state must be specified by enough conditions—such as phase, temperature, pressure, and composition—to make endpoint comparisons meaningful.
Model boundary
Do not infer a unique q or w from ΔE alone without a path or constraint. Temperature is a state variable, but heat is energy transfer and is not a substance stored in a sample.
Misconception
Because ΔE is fixed between two states, every route between them transfers the same heat. Different paths can divide the same ΔE differently between q and w while preserving q + w.
Earlier objectives in this map
System boundaries and the first law
Objective 3GC-THM-LAW-03

Second and third laws

Use ΔS_univ = ΔS_sys + ΔS_surr to evaluate a spontaneous direction and state the third-law reference for absolute entropy.

Must know
A spontaneous process has ΔS_univ > 0, equilibrium has ΔS_univ = 0, and the reverse direction is spontaneous when the proposed direction gives ΔS_univ < 0. The third law assigns zero entropy to a perfect crystal at absolute zero, providing a reference for tabulated absolute molar entropies.
Model boundary
The entropy of the system alone is not a universal spontaneity test. Real substances can retain residual entropy when the perfect-crystal assumptions are not met.
Misconception
A spontaneous process must make the system more disordered. Spontaneity requires a positive total entropy change for system plus surroundings; the system's entropy can decrease when the surroundings increase enough.
Earlier objectives in this map
System boundaries and the first law
BHess’s law3 objectives
Objective 1GC-THM-HES-01

Transform thermochemical equations

Reverse or scale a balanced thermochemical equation and transform its enthalpy change with the reaction exactly as written.

Must know
Reversing a reaction reverses the sign of ΔH because the initial and final states exchange roles. Multiplying every coefficient by a factor multiplies ΔH by the same factor because reaction enthalpy is tied to the stated stoichiometric amount.
Model boundary
Phase labels and conditions are part of a thermochemical equation. Do not cancel, reverse, or scale species without applying the identical operation to the complete reaction and its ΔH.
Misconception
Doubling a reaction leaves ΔH unchanged because enthalpy is a state function. The molar enthalpy for a specified process is intensive, but the ΔH attached to a written reaction scales with its stoichiometric extent.
Earlier objectives in this map
State functions and path functions
Objective 2GC-THM-HES-02

Hess cycle construction

Reverse, scale, and add supplied component reactions so intermediates cancel and the summed equation and enthalpy match a target reaction.

Must know
Because enthalpy is a state function, the ΔH values of component steps add to the ΔH of their exact net reaction. A valid Hess construction must reproduce every target coefficient, species, phase, and direction after intermediates cancel.
Model boundary
Hess's law determines the net enthalpy change, not the mechanism, rate, or likelihood that the listed component reactions occur as elementary steps.
Misconception
Component ΔH values should be averaged when their reactions are combined. After each equation and ΔH are transformed consistently, the component enthalpy changes are added algebraically.
Earlier objectives in this map
Transform thermochemical equations
Objective 3GC-THM-HES-03

Standard enthalpies of formation

Calculate a standard reaction enthalpy from tabulated standard enthalpies of formation using stoichiometric coefficients and explicit species phases.

Must know
ΔH°rxn = ΣνΔHf°(products) − ΣνΔHf°(reactants), with each value multiplied by its balanced coefficient. An element has ΔHf° = 0 only in its reference standard state under the stated conditions; another allotrope or phase does not inherit that value.
Model boundary
Use the tabulated standard state, phase, and temperature as given. A zero formation enthalpy is not a claim that an element contains no internal energy.
Misconception
Every pure element has a standard enthalpy of formation of zero in every phase. Zero applies to the element's reference standard state, not automatically to other allotropes or phases.
Earlier objectives in this map
Hess cycle construction
CSpontaneity3 objectives
Objective 1GC-THM-SPO-01

Spontaneity is not speed

Classify a proposed process as thermodynamically spontaneous, nonspontaneous, or at equilibrium without inferring how rapidly it occurs.

Must know
Spontaneous means thermodynamically favored under the stated conditions; it does not mean instantaneous, irreversible, or complete. Reaction rate and activation barrier belong to kinetics, while spontaneity and equilibrium position are thermodynamic questions.
Model boundary
Do not infer a rate constant, mechanism, or activation energy from ΔG, ΔS_univ, or K alone. A catalyst can change the approach rate without changing the equilibrium condition.
Misconception
A spontaneous reaction must proceed quickly once the reactants are mixed. A thermodynamically favored process can be extremely slow when a large kinetic barrier blocks the path.
Earlier objectives in this map
Second and third laws
Foundation review
Dynamic equilibrium ↗
Objective 2GC-THM-SPO-02

Gibbs-energy direction criterion

Use the sign of ΔG at constant temperature and pressure to determine the favored direction and identify the equilibrium boundary for the current composition.

Must know
At constant temperature and pressure, ΔG < 0 favors the reaction as written, ΔG > 0 favors the reverse direction, and ΔG = 0 marks equilibrium. ΔG describes the current composition and conditions, whereas ΔG° describes a standard-state reaction; they are not interchangeable away from standard conditions.
Model boundary
Use the constant-temperature, constant-pressure criterion and the reaction direction as written. Detailed chemical-potential derivations are outside this objective unless supplied.
Misconception
A negative standard Gibbs energy means every possible mixture has ΔG < 0. Composition changes the reaction Gibbs energy; a system reaches ΔG = 0 at equilibrium even when ΔG° is negative.
Earlier objectives in this map
Spontaneity is not speed
Objective 3GC-THM-SPO-03

Temperature and Gibbs energy

Apply ΔG = ΔH − TΔS with compatible units and temperature in kelvins to predict how enthalpy and entropy signs control spontaneity across temperature regimes.

Must know
When ΔH and ΔS have opposite signs, the process is favored at all temperatures or at none under the constant-sign model; when their signs match, a threshold temperature may separate favored regimes. Convert ΔH and ΔS to compatible energy units and use absolute temperature before subtracting TΔS.
Model boundary
Treat ΔH and ΔS as temperature-independent only when the introductory model or supplied data permit it. A calculated crossover is conditional on phase stability and the stated pressure.
Misconception
The TΔS term can use degrees Celsius because only temperature differences require kelvins. Absolute temperature multiplies entropy in the Gibbs equation, so T must be in kelvins.
Earlier objectives in this map
Gibbs-energy direction criterion
DEnthalpies and entropies3 objectives
Objective 1GC-THM-ESE-01

Enthalpy signs and constant-pressure heat

Interpret the sign and magnitude of a reaction or phase-change enthalpy and relate ΔH to heat transferred at constant pressure.

Must know
At constant pressure with only pressure-volume work, qₚ = ΔH; an exothermic system has ΔH < 0 and an endothermic system has ΔH > 0. Reaction enthalpy depends on direction, stoichiometric amount, phase identity, and conditions, so a sign or value belongs to the process as written.
Model boundary
Do not replace ΔE with ΔH automatically at constant volume or when additional work terms matter. Calorimeter heat and reaction heat have opposite signs when their energy exchange is isolated.
Misconception
An exothermic process has positive ΔH because the surroundings gain heat. For the reacting system, energy leaves as heat at constant pressure, so its enthalpy change is negative.
Earlier objectives in this map
State functions and path functions
Objective 2GC-THM-ESE-02

Entropy and multiplicity trends

Predict the sign of a system entropy change from defensible changes in phase, temperature, gas-particle count, volume, or mixing while naming the conditions behind the comparison.

Must know
Entropy increases with the number of accessible microstates; common positive trends include heating, expansion, mixing distinguishable particles, and moving from solid to liquid to gas under comparable conditions. For a gas-phase reaction under similar conditions, an increase in total gas moles often supports ΔS > 0, but molecular complexity, phase, and data outrank that shortcut.
Model boundary
Qualitative trend rules do not replace tabulated entropy calculations when competing effects are substantial. Avoid treating 'disorder' as a measurable substance or a universal one-word proof.
Misconception
Every reaction that produces more total moles has positive ΔS. Phase and accessible microstates matter; counting total stoichiometric moles without distinguishing gases can give the wrong sign.
Earlier objectives in this map
Second and third laws
Objective 3GC-THM-ESE-03

Standard reaction enthalpy and entropy

Calculate ΔH°rxn and ΔS°rxn from tabulated standard molar quantities, preserving stoichiometric coefficients, phases, product-minus-reactant order, and units.

Must know
For either standard-state property, sum coefficient-weighted product values and subtract coefficient-weighted reactant values. Standard molar entropies are generally not zero for elements at ordinary temperatures, even though reference-state elements have ΔHf° = 0.
Model boundary
Keep J and kJ units explicit before combining results in a Gibbs calculation. Use only the species phases and standard conditions represented by the table.
Misconception
Because an element's standard enthalpy of formation is zero, its standard molar entropy is also zero. The third-law zero is for a perfect crystal at absolute zero; tabulated standard molar entropies at ordinary temperatures are generally positive.
Earlier objectives in this map
Standard enthalpies of formation · Enthalpy signs and constant-pressure heat · Entropy and multiplicity trends
EHeat transfer3 objectives
Objective 1GC-THM-HEA-01

Sensible heat and thermal capacity

Calculate heat transfer from q = mcΔT or q = CΔT, choosing the correct heat-capacity quantity and preserving the sign of the temperature change.

Must know
Specific heat capacity is per unit mass, so q = mcΔT; the heat capacity C of a whole object already includes its amount, so q = CΔT. ΔT = Tfinal − Tinitial has the same numerical increment in Celsius and kelvins, while q is positive for the object whose temperature rises in the isolated sensible-heating model.
Model boundary
Assume a constant heat capacity and no phase change only when the temperature interval and model justify it. Temperature equality at equilibrium does not mean the two objects transferred equal-sign heat.
Misconception
The larger object always undergoes the larger temperature change when two objects exchange heat. Temperature change depends on transferred heat divided by heat capacity; mass alone is insufficient without specific heat.
Earlier objectives in this map
System boundaries and the first law
Objective 2GC-THM-HEA-02

Calorimeter energy balance

Use an isolated calorimeter energy balance to solve for reaction, solution, or calorimeter heat and distinguish constant-pressure from constant-volume measurements.

Must know
For the defined isolated assembly, qrxn + qsurr + qcal = 0, with any negligible term omitted only when the model says so. A coffee-cup calorimeter commonly approximates constant pressure and reports qₚ = ΔH, while a bomb calorimeter operates at constant volume and directly reports qᵥ = ΔE for the modeled reaction.
Model boundary
Use the supplied calorimeter constant, solution mass, density, and heat-capacity assumptions. Do not equate ΔH and ΔE for a gas-producing reaction without a valid conversion or approximation.
Misconception
The reaction and calorimeter have heat values with the same sign because they reach the same final temperature. Heat lost by one part is gained by another, so their contributions to an isolated energy balance have opposite signs.
Earlier objectives in this map
Sensible heat and thermal capacity · State functions and path functions
Objective 3GC-THM-HEA-03

Multistep heating and phase change

Construct and sum signed sensible-heat and phase-change steps for a heating or cooling path that crosses one or more phase boundaries.

Must know
Within one phase, use q = mcΔT with that phase's heat capacity; at an ideal phase-change temperature, use q = nΔHphase while temperature remains constant. Break the full path at every phase boundary, give each step the direction-appropriate sign, and add the step energies algebraically.
Model boundary
Use pressure-specific transition temperatures and enthalpies when supplied. Superheating, supercooling, heat loss, and temperature-dependent heat capacities are outside the ideal segment model unless included.
Misconception
Temperature continues rising while a pure substance melts at its equilibrium melting point. Added energy changes phase while temperature remains constant in the ideal constant-pressure plateau model.
Earlier objectives in this map
Sensible heat and thermal capacity · Enthalpy signs and constant-pressure heat

Free source ladder

Trace every boundary.

Official 2026 DAT scopeDefines the five published topic labels—not their weights. ↗OpenStax · Energy BasicsSystem boundaries, heat, work, internal energy, thermal capacity, and the first-law energy ledger. ↗OpenStax · CalorimetrySensible heat, calorimeter energy balance, constant-pressure and constant-volume models, and phase-change paths. ↗OpenStax · EnthalpyState functions, thermochemical equations, Hess cycles, and standard enthalpies of formation. ↗OpenStax · SpontaneityThermodynamic favorability, direction, and the essential separation between spontaneity and speed. ↗OpenStax · EntropyMicrostates, energy and matter dispersal, phase trends, gas-particle count, and standard molar entropy. ↗OpenStax · The Second and Third LawsTotal entropy, spontaneous direction, equilibrium, and the perfect-crystal reference at absolute zero. ↗OpenStax · Free EnergyGibbs-energy direction, standard versus current conditions, temperature regimes, and equilibrium boundaries. ↗