-818
kJ/mol
-818
kJ/mol
-890.3
kJ/mol
818
kJ/mol
-1
-818
kJ/mol
-818
kJ/mol
-890.3
kJ/mol
818
kJ/mol
-1
The Heat of Reaction Calculator determines the enthalpy change (ΔHrxn) for any chemical reaction using two methods: bond energies and formation enthalpies. The bond energy method estimates ΔH by comparing the energy required to break bonds in reactants with the energy released when forming bonds in products. The formation enthalpy method uses Hess's law with tabulated ΔH°f values. This dual approach is valuable because bond energies give quick estimates for gas-phase reactions while formation enthalpies provide more accurate results for any phase. The calculator is essential for chemical kinetics, process engineering, and reaction feasibility assessment.
Method 1 — Bond Energies:
$$\Delta H_{rxn} = \sum D(\text{bonds broken}) - \sum D(\text{bonds formed})$$
Breaking bonds requires energy (endothermic, positive), forming bonds releases energy (exothermic, negative). If more energy is released in forming bonds than consumed in breaking bonds, ΔH < 0 (exothermic).
Method 2 — Formation Enthalpies:
$$\Delta H^\circ_{rxn} = \sum \Delta H^\circ_f(\text{products}) - \sum \Delta H^\circ_f(\text{reactants})$$
This method is more accurate because it accounts for molecular interactions beyond simple bond energies, including resonance, strain, and intermolecular forces.
A negative ΔHrxn means the reaction is exothermic — it releases energy and the products are more stable. A positive ΔHrxn means endothermic — energy input is required. The energy difference shows the absolute magnitude regardless of direction. Bond energy estimates are typically accurate to ±10–20 kJ/mol, while formation enthalpies are accurate to ±1 kJ/mol.
Inputs
Results
Bonds broken: 4(C−H) + 2(O=O) = 4(413) + 2(498) = 2648 kJ. Bonds formed: 2(C=O) + 4(O−H) = 2(799) + 4(467) = 3466 kJ. ΔH = 2648 − 3466 = −818 kJ (exothermic, approximate).
Inputs
Results
ΔH° = [−393.5 + 2(−285.8)] − [−74.8 + 0] = −965.1 + 74.8 = −890.3 kJ. More accurate than bond energy estimate (−818 vs −890.3 kJ).
Bond energies are averages over many molecules, ignoring specific molecular environments, resonance, and intermolecular forces. Formation enthalpies are specific to exact compounds and phases, making them more accurate.
Use bond energies for quick estimates, especially for gas-phase organic reactions. They are useful when ΔH°f values are unavailable or for understanding which bonds drive the energetics.
Bond dissociation energy (D) is the energy required to break one mole of a specific bond in a specific molecule. Average bond energies used in this calculator are averages across many different molecules.
ΔH alone does not determine feasibility. You need ΔG = ΔH − TΔS. However, strongly exothermic reactions (large negative ΔH) are often feasible because the enthalpy term dominates.
C−H: 413 kJ/mol, C−C: 348, C=C: 614, C≡C: 839, O=O: 498, N≡N: 941, C=O: 799, O−H: 467, H−H: 436 kJ/mol. These are average values.
Bond energies are not appropriate for ionic reactions (no covalent bonds to break/form). Use formation enthalpies instead, which account for ionic bonding, lattice energy, and solvation.
For molecules with resonance (like benzene), use the actual bond energy (not hypothetical single/double bond energies). The resonance stabilization energy is the difference between calculated and actual values.
ΔHrxn is the overall energy change, not the activation energy (Ea). A reaction can be exothermic but still have a high activation barrier, requiring a catalyst or heating to proceed.
Yes, if the energy of bonds broken exactly equals bonds formed (or product and reactant formation enthalpies are equal). This is rare but possible for thermoneutral reactions.
ΔH varies with temperature via Kirchhoff's equation: ΔH(T₂) = ΔH(T₁) + ∫ΔCp dT. For moderate temperature ranges, ΔH is approximately constant.
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