Exothermic vs Endothermic Predictor

Predict exothermic or endothermic nature from bond energies, formation enthalpies, or Gibbs free energy.

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Where Reaction Enthalpy Comes From

Breaking bonds requires energy; forming bonds releases it. The overall enthalpy change is simply the balance. If the bonds formed are stronger than those broken, energy is released and the reaction is exothermic.

ΔH°rxn = ΣΔH°f(products) − ΣΔH°f(reactants)

Standard enthalpies of formation are defined relative to elements in their standard states, which are assigned zero by convention. This is why ΔH°f of O2, N2 and graphite are all zero — not because they contain no energy, but because they are the reference point.

Sign of ΔHNameEnergy flowFeelsExamples
NegativeExothermicReleased to surroundingsWarmCombustion, neutralisation, most oxidations
PositiveEndothermicAbsorbed from surroundingsColdPhotosynthesis, thermal decomposition, most dissolving

Exothermic Is Not the Same as Spontaneous

This is the distinction that matters most and is most often missed. Spontaneity depends on ΔG, not ΔH:

ΔG = ΔH − TΔS

Ammonium nitrate dissolving in water is strongly endothermic — the solution gets noticeably cold, which is how instant cold packs work — yet it happens spontaneously. The large positive entropy change of the ordered crystal dispersing into solution outweighs the unfavourable enthalpy term.

The reverse case also exists. Many exothermic reactions with strongly negative ΔS become non-spontaneous above a crossover temperature, which is why ammonia synthesis yields fall as the Haber process is heated.

Hess’s Law

Because enthalpy is a state function, ΔH depends only on initial and final states, not the route taken. This allows enthalpies to be added, reversed (changing sign) and scaled — which is how values for reactions too difficult to measure directly are obtained from ones that can be.

Worked Examples

Example 1: Combustion CH4: Hf products=-393.5+2x(-285.8)=-965.1, reactants=-74.8
dH=-965.1-(-74.8)
Result: dH=-890.3 kJ/mol EXOTHERMIC
Natural gas combustion
Example 2: Photosynthesis (reverse of combustion)
dH=+890.3 kJ/mol
Result: ENDOTHERMIC - driven by light energy
Plants absorb energy from sunlight
Example 3: Spontaneous but endothermic
NH4NO3 dissolving: ΔH = +25.7 kJ/mol, ΔS = +108 J/(mol·K)
Result: ΔG = −6.5 kJ/mol at 298 K — spontaneous
The entropy term TΔS = 32.2 kJ/mol outweighs the endothermic enthalpy. This is exactly how instant cold packs work.
Example 4: Using Hess’s law
Target reaction unmeasurable directly, but two known reactions sum to it
Result: Add their ΔH values
Reversing a reaction flips the sign of ΔH; doubling it doubles ΔH. Enthalpy is a state function, so any valid route gives the same answer.
Example 5: Bond energy estimate
Bonds broken 2,650 kJ, bonds formed 3,540 kJ
Result: ΔH ≈ −890 kJ/mol
Close to the tabulated value for methane combustion. Bond energy estimates are approximate because tabulated values are averages across many molecules.

Common Mistakes

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Equating exothermic with spontaneous

Spontaneity is governed by ΔG. Ammonium nitrate dissolves spontaneously while absorbing heat, because the entropy increase dominates.

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Reversing the subtraction order

It is products minus reactants. Reversing gives the right magnitude with the wrong sign, turning exothermic into endothermic.

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Forgetting stoichiometric coefficients

Each ΔH°f is multiplied by its coefficient in the balanced equation. Two moles of water contributes twice its formation enthalpy.

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Assuming elements always have ΔH°f = 0

Only in their standard states. Ozone is not zero, and neither is monatomic oxygen — only O2 gas at standard conditions.

Frequently Asked Questions

Spontaneous vs exothermic?
Many exothermic reactions are spontaneous at room temperature. But some endothermic reactions are also spontaneous (ice melting: endothermic but entropy gain drives it). Always check Gibbs free energy for spontaneity.
Temperature dependence?
Exothermic + entropy decrease: spontaneous only below certain T. Endothermic + entropy increase: spontaneous only above certain T. Calculate T_crossover = delta-H/delta-S.
What is the difference between exothermic and spontaneous?
Exothermic means heat is released, determined by ΔH. Spontaneous means the reaction proceeds without input, determined by ΔG. Endothermic reactions can be spontaneous when entropy increases enough.
Why do elements have zero enthalpy of formation?
By convention, elements in their standard states are the reference point of the scale. It is a chosen zero, not a statement that they contain no energy.
How does an instant cold pack work?
Ammonium nitrate dissolving is endothermic, absorbing about 25.7 kJ/mol from the surroundings. It happens spontaneously because the entropy gain outweighs the enthalpy cost.
What is Hess’s law?
Enthalpy change depends only on initial and final states, not the path. Reactions can therefore be added, reversed or scaled to find ΔH for reactions that cannot be measured directly.
Can I calculate ΔH from bond energies?
Approximately — bonds broken minus bonds formed. Tabulated bond energies are averages across many compounds, so results typically differ from formation-enthalpy calculations by a few percent.

Formula Explorer connections

Interpretation: This relationship tracks energy transfer, state-function change or the balance between enthalpy and entropy in a chemical process. Assumption: Keep energy units compatible, use kelvin for absolute temperature, and match standard states and reaction stoichiometry. Thermodynamic favorability does not determine reaction speed.

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