Heat of Combustion Calculator
Calculate the molar heat of combustion of any fuel using either experimental calorimetry data or standard formation enthalpies. Select a fuel preset for instant look-up values, or enter your own mass, molar mass and calorimeter readings. Results update in kJ/mol, kJ/g and MJ/kg as you type, with a full step-by-step worked solution showing every formula.
Formula
Worked example
Calorimetry example: 0.80 g of methane burned, raising 200 g of water by 9.5 °C. Q = 200 x 4.184 x 9.5 = 7949 J = 7.949 kJ. Moles burned = 0.80 / 16.04 = 0.04990 mol. Heat of combustion = 7.949 / 0.04990 = 159.3 kJ/mol (a student experiment value; the tabulated value is 890 kJ/mol due to heat losses in open apparatus).
What is the heat of combustion?
The heat of combustion (also called molar enthalpy of combustion, deltaHc) is the energy released as heat when one mole of a substance undergoes complete combustion with oxygen under standard conditions (25 °C, 1 atm). By convention, the value is reported as a positive number representing the exothermic magnitude, even though the corresponding reaction enthalpy is negative. For example, burning one mole of methane releases 890 kJ: CH4 + 2 O2 → CO2 + 2 H2O, deltaH = -890 kJ/mol. The same quantity expressed per gram is 890/16.04 = 55.5 kJ/g, and per kilogram it is 55.5 MJ/kg. These per-mass values are the standard way engineers compare fuel energy densities.
Two ways to calculate it
You can measure the heat of combustion directly using a calorimeter, or calculate it indirectly from standard formation enthalpies using Hess's law. Calorimetry method: a known mass of fuel is burned inside a calorimeter containing a known mass of water. The temperature rise of the water gives Q = m x c x deltaT (where c for water = 4.184 J/g°C). Dividing Q by the moles of fuel burned gives the molar heat of combustion. Coffee-cup calorimeters (open, constant pressure) are common in teaching labs; bomb calorimeters (sealed, constant volume) give more accurate results because no heat escapes. Formation enthalpy method (Hess's law): deltaH_comb = sum of formation enthalpies of products minus the formation enthalpy of the fuel. Because the reaction is exothermic, this sum is negative; the magnitude is the heat of combustion. This approach works without any experiment, using only tabulated standard formation enthalpies.
Higher and lower heating values
For fuels that contain hydrogen, the combustion products include water. Whether that water is liquid or vapour matters because condensing steam releases additional heat (the enthalpy of vaporisation, about 44 kJ/mol of water at 25 °C). The higher heating value (HHV) assumes water condenses to liquid and is used in most chemistry tables. The lower heating value (LHV) assumes water remains as vapour and is typically used in engineering efficiency calculations. For methane, HHV is 890 kJ/mol and LHV is about 802 kJ/mol. This calculator uses HHV (liquid water convention) for the reference table and for the Hess's law mode when standard liquid-water formation enthalpies are used.
Sources of error in calorimetry experiments
Student calorimeter results almost always fall below the tabulated value. Common reasons include: incomplete combustion (yellow soot instead of full oxidation to CO2 and water); heat loss to the surrounding air, calorimeter walls and stand; the specific heat of the calorimeter apparatus itself is ignored; evaporation of water changes the effective mass. A well-designed bomb calorimeter corrects for apparatus heat capacity and uses oxygen at high pressure to ensure complete combustion, bringing results within 0.1% of tabulated values.
Standard heats of combustion for common fuels
| Fuel | Formula | Molar mass (g/mol) | deltaHc (kJ/mol) | Heating value (MJ/kg) |
|---|---|---|---|---|
| Hydrogen | H2 | 2.016 | 286 | 141.8 |
| Carbon | C | 12.011 | 394 | 32.8 |
| Methane | CH4 | 16.04 | 890 | 55.5 |
| Ethane | C2H6 | 30.07 | 1560 | 51.9 |
| Propane | C3H8 | 44.10 | 2220 | 50.3 |
| Butane | C4H10 | 58.12 | 2877 | 49.5 |
| Pentane | C5H12 | 72.15 | 3537 | 49.0 |
| Hexane | C6H14 | 86.18 | 4163 | 48.3 |
| Octane | C8H18 | 114.23 | 5460 | 47.8 |
| Methanol | CH3OH | 32.04 | 726 | 22.7 |
| Ethanol | C2H5OH | 46.07 | 1368 | 29.7 |
| Glucose | C6H12O6 | 180.16 | 2803 | 15.6 |
Values are the positive (exothermic) magnitudes at standard conditions (25 °C, 1 atm), liquid water as product (higher heating value basis). Source: NIST / LibreChem.
Frequently asked questions
Why is heat of combustion always reported as a positive number?
By thermochemical convention, the molar enthalpy change for combustion is negative (deltaH < 0) because the reaction releases energy. However, "heat of combustion" is defined as the magnitude of that enthalpy change, so it is quoted as a positive number for clarity in engineering and educational contexts. When you see deltaHc = 890 kJ/mol for methane, that means 890 kJ is released; the reaction enthalpy is -890 kJ/mol.
What is the difference between heat of combustion and heating value?
These terms mean the same thing but are used in different contexts. Chemists prefer "molar heat (or enthalpy) of combustion" in kJ/mol to compare molecules. Engineers prefer "heating value" in MJ/kg (or MJ/m3 for gases) to compare fuels by mass or volume. Numerically, kJ/g and MJ/kg are identical (both equal to kJ/mol divided by molar mass in g/mol).
What is the formula for heat of combustion from calorimetry?
Three steps: (1) Q = m x c x deltaT, where m is the mass of water in grams, c is the specific heat capacity of water (4.184 J/g°C), and deltaT is the temperature rise in °C. This gives the heat absorbed in joules; divide by 1000 to get kJ. (2) n = m_fuel / M, where m_fuel is the mass of fuel burned and M is its molar mass. (3) deltaHc = Q / n, giving kJ per mole of fuel.
How do I use Hess's law to find heat of combustion?
Write the balanced combustion equation for your fuel. Look up the standard enthalpy of formation for each product (CO2, H2O, SO2 if sulfur is present) and multiply each by its stoichiometric coefficient. Subtract the formation enthalpy of the fuel. The result is the reaction enthalpy (negative for exothermic reactions). The heat of combustion is the absolute value. For methane: products give 1 x (-393.5) + 2 x (-285.8) = -965.1 kJ/mol; minus fuel formation of -74.8 kJ/mol gives -965.1 - (-74.8) = -890.3 kJ/mol, so deltaHc = 890 kJ/mol.
Which fuel has the highest heat of combustion per kilogram?
Hydrogen has by far the highest gravimetric heating value at about 141.8 MJ/kg, more than twice that of any hydrocarbon. Among common hydrocarbons, methane is highest at 55.5 MJ/kg, slightly ahead of other alkanes (propane 50.3, octane 47.8). Alcohols are lower because they already contain oxygen: ethanol delivers about 29.7 MJ/kg. Glucose is lowest among fuels at about 15.6 MJ/kg.
Why does my lab result differ so much from the textbook value?
Most school calorimetry experiments use open spirit burners or simple metal calorimeters, which lose a large fraction of heat to the surroundings and suffer from incomplete combustion. A result 20-50% below the tabulated value is typical. To improve accuracy: use a draught shield around the flame, keep the water container close to the burner, weigh the fuel before and after with a balance precise to 0.001 g, and record temperatures to 0.1 °C. Calculating the percentage error against the tabulated value is a standard part of the exercise.