ChemistryDifficulty 3-5

Reaction Energy

Every chemical reaction involves an energy change, because breaking bonds absorbs energy and forming bonds releases it. If more energy is released than absorbed, the reaction is exothermic (ΔH < 0) and warms its surroundings; if the reverse, it is endothermic (ΔH > 0). Separately, almost every reaction must first climb an energy barrier, the activation energy, which sets how fast it goes — and which a catalyst lowers without being used up.

Thermochemistry: Hess's law as a consistency check

Because enthalpy is a state function, ΔH for an overall process equals the sum over any sequence of steps. Formation enthalpies therefore let us compute reaction enthalpies we never measured directly, and comparing with calorimetry tests the data: the formation-enthalpy route gives −890.6 kJ/mol for methane combustion, the direct measurement −890.7 ± 0.4 kJ/mol.

Kinetics: energy profile

On a reaction energy profile, the transition state is the maximum between reactants and products; Eₐ is its height above the reactants, and ΔH is the difference between reactant and product levels. The Arrhenius equation gives the temperature dependence; a plot of ln k against 1/T is a straight line of slope −Eₐ/R.

\ln k = \ln A - \frac{E_a}{R}\cdot\frac{1}{T}

Common misconception: ΔH alone does not determine Eₐ: reactions with similar ΔH can have very different barriers. On an energy profile the forward and reverse barriers differ by ΔH, so an endothermic step's barrier is at least as large as its ΔH.

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Where reaction energy comes from

Energy is conserved: it is never created or destroyed, only transformed, and it is measured in joules. In a reaction, breaking the reactants' bonds costs energy and forming the products' bonds releases energy. The difference appears as heat exchanged with the surroundings, described by the enthalpy change ΔH.

Exothermic (ΔH < 0): heat flows out to the surroundings — burning methane is an example.
Endothermic (ΔH > 0): heat is absorbed from the surroundings — splitting water into hydrogen and oxygen is an example.

Worked example: burning methane

\Delta H^\circ_{\text{rxn}} = \sum m\,\Delta_f H^\circ(\text{products}) - \sum n\,\Delta_f H^\circ(\text{reactants})

Standard enthalpy of reaction from standard enthalpies of formation (elements in their standard states are zero).

For CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l), using NIST values ΔfH°(CO₂, g) = −393.51, ΔfH°(H₂O, l) = −285.83 and ΔfH°(CH₄, g) = −74.6 kJ/mol, and zero for O₂: ΔH° = [−393.51 + 2(−285.83)] − [−74.6] ≈ −890.6 kJ/mol (ScienceVerse calculation). That agrees with the directly measured enthalpy of combustion, −890.7 ± 0.4 kJ/mol — a check that Hess's law works.

Reversing a reaction reverses the sign of ΔH. Because forming liquid water from its elements releases 285.83 kJ/mol, splitting liquid water back into hydrogen and oxygen requires +285.83 kJ/mol (ScienceVerse calculation).

Common misconception: A common belief is that energy is 'stored in' bonds and released when they break. Breaking any bond absorbs energy; fuels release energy because the bonds formed in the products (here in CO₂ and H₂O) are stronger overall than those broken.

The energy barrier: activation energy

Even an exothermic reaction usually needs a push to start. Colliding particles must pass through a transition state — the highest-energy arrangement on the way to products — and the energy needed to reach it is the activation energy, Eₐ.

k = A\,e^{-E_a/RT}

Arrhenius equation: the rate constant k rises with temperature T and falls with activation energy Eₐ.

Common misconception: ΔH and Eₐ answer different questions. ΔH says how much heat a reaction gives out or takes in overall; Eₐ governs how fast it goes. A strongly exothermic reaction can still be extremely slow if its activation energy is high.

Catalysts

A catalyst speeds a reaction by offering a pathway with a lower activation energy, and is regenerated rather than consumed. Wilhelm Ostwald showed this in 1894 and received the 1909 Nobel Prize in Chemistry partly for his work on catalysis. Catalysts are used to make sulfuric acid, ammonia and methanol on an industrial scale, and enzymes — biological catalysts — speed reactions in cells by factors often between a million and 10²⁰; without them most of those reactions would be too slow to sustain life.

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Sources and methodology

A negative enthalpy change (ΔH < 0) indicates an exothermic reaction and a positive one (ΔH > 0) an endothermic reaction, and if the direction of a chemical equation is reversed, the sign of its ΔH changes. (awaiting scientific review)
- 5.4: Enthalpy (Chemistry 2e, OpenStax, on LibreTexts), CC BY 4.0 — Textbook
A chemical or physical change that releases heat is called an exothermic process, and one that absorbs heat is an endothermic process. (awaiting scientific review)
- 5.2: Energy Basics (Chemistry 2e, OpenStax, on LibreTexts), CC BY 4.0 — Textbook
During a chemical reaction, breaking bonds in the reactants requires energy and forming new bonds in the products releases energy; bond energies, defined for breaking one mole of a bond in gaseous molecules, are therefore always positive. (awaiting scientific review)
- 5.8: Bond Energies (Chemistry: Physical Principles), CC BY 4.0 — Textbook
Energy is conserved — it cannot be created or destroyed, only transformed from one form to another — and its SI unit is the joule. (awaiting scientific review)
- 5.2: Energy and Its Conversion (Chemistry: Physical Principles), CC BY 4.0 — Textbook
Hess's law states that if a process can be written as the sum of several steps, the enthalpy change of the whole process equals the sum of the enthalpy changes of the steps; it holds because enthalpy is a state function, depending only on where a process starts and ends. (awaiting scientific review)
- 5.4: Enthalpy (Chemistry 2e, OpenStax, on LibreTexts), CC BY 4.0 — Textbook
A standard enthalpy of reaction can be calculated as the sum of the standard enthalpies of formation of the products minus the sum for the reactants, each multiplied by its stoichiometric coefficient; the standard enthalpy of formation of an element in its most stable form is zero. (awaiting scientific review)
- 5.4: Enthalpy (Chemistry 2e, OpenStax, on LibreTexts), CC BY 4.0 — Textbook
The standard enthalpy of combustion of methane gas has been measured as −890.7 ± 0.4 kJ/mol. (awaiting scientific review)
- NIST Chemistry WebBook (SRD 69): Methane — gas phase thermochemistry data — Primary dataset
The standard enthalpy of formation of methane gas is −74.6 ± 0.3 kJ/mol (Manion's 2002 review value, as listed by NIST). (awaiting scientific review)
- NIST Chemistry WebBook (SRD 69): Methane — gas phase thermochemistry data — Primary dataset
The standard enthalpy of formation of liquid water is −285.830 ± 0.040 kJ/mol (CODATA review value). (awaiting scientific review)
- NIST Chemistry WebBook (SRD 69): Water — condensed phase thermochemistry data — Primary dataset
The standard enthalpy of formation of gaseous carbon dioxide is −393.51 ± 0.13 kJ/mol (CODATA review value). (awaiting scientific review)
- NIST Chemistry WebBook (SRD 69): Carbon dioxide — gas phase thermochemistry data — Primary dataset
The transition state is the highest-energy arrangement between reactants and products, so the activation energy is the energy barrier that colliding particles must overcome for products to form. (awaiting scientific review)
- 6.6: Collision Theory, Activation Energy, and the Arrhenius Equation (Chemistry: Physical Principles), CC BY 4.0 — Textbook
The Arrhenius equation, k = A·exp(−Eₐ/RT), shows that a higher temperature and a lower activation energy both give a larger rate constant. (awaiting scientific review)
- 6.2.3.1: Arrhenius Equation (Stephen Lower), CC BY — Textbook
Plotting ln k against 1/T gives a straight line whose slope is −Eₐ/R, so the activation energy can be found from rate constants measured at different temperatures. (awaiting scientific review)
- 6.6: Collision Theory, Activation Energy, and the Arrhenius Equation (Chemistry: Physical Principles), CC BY 4.0 — Textbook
On a reaction energy profile the activation energies of the forward and reverse reactions differ by the enthalpy change: for an exothermic reaction the reverse barrier is larger than the forward one by an amount equal to ΔH. (awaiting scientific review)
- 6.6: Collision Theory, Activation Energy, and the Arrhenius Equation (Chemistry: Physical Principles), CC BY 4.0 — Textbook
A catalyst increases a reaction's rate by providing a pathway with a lower activation energy; it lowers the barrier for both the forward and reverse reactions and does not change the equilibrium constant. (awaiting scientific review)
- 6.8: Catalysis (Chemistry: Physical Principles), CC BY 4.0 — Textbook
Wilhelm Ostwald, awarded the 1909 Nobel Prize in Chemistry in recognition of his work on catalysis, showed in 1894 that a catalyst affects the speed of a chemical reaction but is not part of its end-products. (awaiting scientific review)
- The Nobel Prize in Chemistry 1909: Wilhelm Ostwald — Facts — Other (unclassified)
Important industrial processes such as the preparation of sulfuric acid, ammonia and methanol use heterogeneous catalysts, and enzymes, the biological catalysts, increase reaction rates often by factors of 10⁶ to 10²⁰ — without catalysts most reactions in cells would be too slow to sustain life. (awaiting scientific review)
- 6.8: Catalysis (Chemistry: Physical Principles), CC BY 4.0 — Textbook

Claims marked “awaiting scientific review” cite the sources listed but have not yet been signed off by a scientific reviewer.

Content status: published 1 October 2026.

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