3.5
°C
29,288
J
2,975
J
32,263
J
-32,263
J
-32.263
kJ/g
0.062344
mol
-517.5
kJ/mol
3.5
°C
29,288
J
2,975
J
32,263
J
-32,263
J
-32.263
kJ/g
0.062344
mol
-517.5
kJ/mol
The Heat of Combustion Calculator determines the enthalpy of combustion (ΔHcomb) from bomb calorimetry data. Combustion reactions are complete oxidation reactions that produce CO₂ and H₂O, releasing energy as heat. This calculator processes calorimeter data — including the temperature rise, water mass, and calorimeter heat capacity — to determine the total heat released and convert it to kJ per mole of substance burned. The heat of combustion is fundamental for determining fuel energy content, food calorie values, standard enthalpies of formation, and evaluating biofuel efficiency.
In bomb calorimetry, the sample burns at constant volume. The total heat released is absorbed by the water and calorimeter:
$$q_{total} = q_{water} + q_{calorimeter}$$
$$q_{water} = m_{water} \cdot c_{water} \cdot \Delta T$$
$$q_{calorimeter} = C_{cal} \cdot \Delta T$$
where Ccal is the calorimeter's heat capacity (J/°C). The heat of combustion per mole is:
$$\Delta H_{comb} = -\frac{q_{total}}{n} = -\frac{q_{total} \cdot M}{m_{sample}}$$
The negative sign indicates combustion is exothermic. Note: bomb calorimetry measures ΔU (constant volume), and ΔH = ΔU + ΔnRT for the conversion, but for condensed-phase samples this difference is small.
The ΔHcomb is always negative (exothermic). More negative values indicate higher energy content. For comparison: methane ≈ −890 kJ/mol, ethanol ≈ −1367 kJ/mol, glucose ≈ −2803 kJ/mol. The heat per gram is useful for comparing fuels of different molecular weights. The calorimeter contribution often represents 5–15% of the total heat absorption.
Inputs
Results
ΔT = 3.5°C. q_water = 2000 × 4.184 × 3.5 = 29,288 J. q_cal = 850 × 3.5 = 2,975 J. Total = 32,263 J. Moles = 0.5/16.04 = 0.03118. ΔH = −32263/0.03118/1000 = −1034.9 kJ/mol.
Inputs
Results
ΔT = 4.68°C. q_water = 2500 × 4.184 × 4.68 = 48,953 J. q_cal = 1250 × 4.68 = 5,850 J. Total = 54,803 J. Moles = 1.025/128.17 = 0.007997. ΔH_comb ≈ −6853 kJ/mol.
The heat of combustion (ΔHcomb) is the enthalpy change when one mole of a substance completely burns in oxygen. It is always negative because combustion releases energy.
Bomb calorimetry is an experimental technique where a sample burns in a sealed, pressurized container (the 'bomb') surrounded by water. The temperature rise of water and bomb measures the heat released.
Combustion is an exothermic oxidation reaction. The products (CO₂ and H₂O) are more thermodynamically stable than the reactants, so energy is released.
The calorimeter heat capacity (Ccal) accounts for heat absorbed by the bomb, stirrer, thermometer, and container walls. It is determined by calibration with a substance of known ΔHcomb (like benzoic acid).
Food calories (kcal) are determined by bomb calorimetry. 1 food Calorie = 1 kcal = 4184 J. The heat of combustion of proteins, carbohydrates, and fats determines their caloric content.
Bomb calorimetry measures ΔU (constant volume). ΔH = ΔU + ΔnRT, where Δn is the change in moles of gas. For reactions with Δn = 0 or small, ΔH ≈ ΔU.
Modern bomb calorimeters achieve accuracy of ±0.01–0.1%, making them among the most precise thermochemical instruments available.
Bomb calorimetry works for combustible substances: organic compounds, fuels, foods. It does not work for inorganic salts, metals (that do not burn), or substances that react incompletely with oxygen.
The standard heat of combustion (ΔH°comb) is measured at 25°C and 1 atm, with products being CO₂(g) and H₂O(l). These standardized values are tabulated for reference.
Using Hess's law: ΔH°f = ΣΔH°f(products of combustion) − ΔH°comb. This allows calculation of formation enthalpies from combustion data.
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