Hess's Law: Path-Independent Enthalpy
Interactive enthalpy diagrams, a mountain analogy, the carbon monoxide formation cycle, and guided practice for reversing, scaling, and adding thermochemical equations.
About this simulation
- What
- An interactive Chemistry simulation of Hess's Law: Path-Independent Enthalpy.
- Who
- Designed for AP, IB, and A‑Level Chemistry students.
- How
- Runs in any modern browser — drag, adjust, and explore in real time.
Updated 2026-03-20
Key Concepts
State function
Enthalpy H depends only on the state of the system. For the same initial and final states, ΔH is the same no matter how many steps you take.
Constant pressure vs volume
At constant pressure (only expansion work), measured heat equals ΔH. At constant volume, measured heat equals ΔU; the same path-independence idea applies to ΔU.
Thermochemical algebra
Reversing an equation flips the sign of ΔH. Multiplying coefficients multiplies ΔH. Adding equations adds their ΔH values—always keep stoichiometry consistent.
Understanding Hess's Law
Hess's law states that, for a given overall reaction, the standard enthalpy change is the sum of the enthalpy changes of any sequence of steps that connect the same initial and final states—provided each step is written for the same temperature and pressure convention.
This follows because enthalpy is a state function. Laboratory measurements usually target constant-pressure calorimetry, where heat equals ΔH. At constant volume, heat equals ΔU, and the same logical structure applies to ΔU.
The graphite–CO–CO₂ cycle on this page uses standard formation-related data to show how an awkward step (C + ½O₂ → CO) can be obtained from better-characterized combustions.
Frequently Asked Questions
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