Thermochemistry & Hess's Law

What is Thermochemistry & Hess's Law?

Thermochemistry is the branch of chemistry that quantifies energy changes — primarily enthalpy (H) — during chemical reactions at constant pressure. An exothermic reaction releases energy (ΔH < 0): the products sit lower on the energy diagram than the reactants, and the surroundings warm up. An endothermic reaction absorbs energy (ΔH > 0): products sit higher, and the surroundings cool down. ΔH can be calculated from standard enthalpies of formation (ΔH_rxn = ΣΔH_f°(products) − ΣΔH_f°(reactants)), from bond energies, or assembled via Hess's Law — any multi-step pathway whose steps sum to the target equation. This simulation animates energy diagrams for combustion, neutralization, dissolution, and photosynthesis, and lets you build Hess's Law pathways step by step.

Parameters explained

Enthalpy Change(kJ/mol)

Enthalpy Change sets ΔH, the heat absorbed or released by the reaction at constant pressure, in kJ/mol. Negative values model exothermic processes whose products sit lower in enthalpy than the reactants; positive values model endothermic processes that require energy input. In the ΔG equation, ΔH is the starting energy term before entropy and temperature are considered. Move this slider while holding ΔS and T fixed to see that a more negative ΔH pushes ΔG downward, while a more positive ΔH can make spontaneity depend on the TΔS contribution.

Entropy Change(J/(mol·K))

Entropy Change sets ΔS, the change in dispersal or disorder, in J/(mol·K). Positive ΔS means the products have more available microstates than the reactants, such as when gas particles increase or a solid dissolves into mobile ions. Negative ΔS means the products are more ordered. Because the simulation calculates TΔS in kJ/mol, the entropy value is divided by 1000 before combining with ΔH. Keep ΔH fixed and raise temperature to see why positive ΔS becomes increasingly favorable at high T, while negative ΔS becomes a larger penalty.

Temperature(K)

Temperature sets the absolute temperature in kelvin, the scale required for thermodynamic equations. It controls how strongly entropy affects spontaneity through the TΔS term in ΔG = ΔH − TΔS. At low temperature, ΔH often dominates because TΔS is small. At high temperature, entropy can outweigh enthalpy, making some endothermic but disorder-increasing reactions spontaneous. Try the Endothermic (high T) preset, then lower Temperature until ΔG changes sign. This shows that spontaneity is not fixed by ΔH alone; it can depend on classroom or lab conditions.

Common misconceptions

• MisconceptionExothermic reactions are always spontaneous.

  CorrectA negative ΔH helps make ΔG negative, but spontaneity depends on both enthalpy and entropy: ΔG = ΔH − TΔS. If ΔS is strongly negative, the TΔS term becomes a penalty, especially at high temperature. Use a negative ΔH with a negative ΔS to see that an exothermic process can become nonspontaneous under some conditions.

• MisconceptionEndothermic reactions can never be spontaneous because they absorb heat.

  CorrectEndothermic means ΔH is positive, so enthalpy alone is unfavorable. A reaction can still be spontaneous if ΔS is positive and the temperature is high enough for TΔS to exceed ΔH. The Endothermic (high T) preset demonstrates this: the reaction absorbs heat, but the entropy gain can still make ΔG negative.

• MisconceptionEntropy is the same thing as heat.

  CorrectEntropy measures energy dispersal and the number of accessible microstates, not heat itself. Heat flow is connected to enthalpy and temperature, while entropy describes how spread out matter and energy can become. The units are different too: ΔH is kJ/mol, while ΔS is J/(mol·K), so the simulation converts TΔS before comparing it with ΔH.

• MisconceptionTemperature only changes reaction speed, not thermodynamic favorability.

  CorrectTemperature can affect both rate and spontaneity, but in different ways. Kinetics uses temperature to explain how often molecules overcome activation barriers. Thermodynamics uses temperature in ΔG = ΔH − TΔS. Changing Temperature in this simulation changes the TΔS term and can flip whether the reaction is thermodynamically favored.

• MisconceptionA negative ΔG means products form instantly and completely.

  CorrectNegative ΔG means the forward direction is thermodynamically favored under the selected conditions, not that it is fast or goes to 100% completion. A reaction can be spontaneous but slow if the activation barrier is high. The equilibrium readout connects ΔG to favorability, while kinetics would be needed to predict how quickly the change occurs.

How teachers use this lab

1. ΔG sign prediction: choose the Exothermic Spontaneous preset, hide the readout briefly, and ask students to predict the sign of ΔG from ΔH, ΔS, and T before revealing the calculation.
2. Temperature threshold investigation: start with the Endothermic (high T) preset, lower Temperature in steps, and have students identify where ΔG changes sign. Connect the threshold to AP 5.C.1 reasoning about ΔG = ΔH − TΔS.
3. Units and conversion check: have students calculate TΔS by hand using ΔS in J/(mol·K), convert to kJ/mol, and compare their result with the simulation display.
4. Entropy misconception probe: keep ΔH fixed, switch ΔS from positive to negative, and ask students to explain the molecular order changes shown by the particle visualization.
5. Preset comparison CER: assign each group one preset and require a claim about spontaneity, evidence from the three slider values, and reasoning that references both enthalpy and entropy.