The reason fire is hot is because the energy released during the combustion reaction is greater than the activation energy needed to start the reaction. In other words, breaking chemical bonds in the fuel always releases more energy (heat) than is absorbed forming new chemical bonds to make products.
In a typical combustion reaction, fuel combined with oxygen and energy to produce carbon dioxide, water, and more energy.
Fuel + Oxygen + Energy → Carbon Dioxide + Water + More Energy
The starting energy could be a lit match or a spark. Flames and fire are visible evidence of energy release. Flames mostly consist of hot and ionized gases. However, combustion may release heat slowly and invisibly, too.
Why Fire Is Hot
In a nutshell: Fire is hot because it releases more energy than it needs to start the chemical reaction. The excess energy maintains the reaction, converting stored energy in fuel into heat and light.
How Hot Is Fire?
There is no one-size-fits-all temperature for fire. Flame temperature depends on several factors, including the chemical composition of the fuel, atmospheric pressure, oxygen content, and portion of the flame being measured.
Here are typical temperatures of flames produced by various fuels:
| Fuel | Flame Temperature |
| Charcoal | 750–1,200 °C (1,382–2,192 °F) |
| Methane (natural gas) | 900–1,500 °C (1,652–2,732 °F) |
| Kerosene | 990 °C (1,814 °F) |
| Gasoline | 1,026 °C (1,878.8 °F) |
| Wood | 1,027 °C (1,880.6 °F) |
| Candle wax | 1,100 °C (2,012 °F) up to 1,300–1,400 °C (2,372–2,552 °F) |
| Methanol | 1,200 °C (2,192 °F) |
| Propane | 1,200–1,700 °C (2,192–3,092 °F) |
| Charcoal | 1,390 °C (2,534 °F) |
| Magnesium | 1,900–2,300 °C (3,452–4,172 °F) |
| MAPP torch | 2,020 °C (3,668 °F) |
| Acetylene torch | Up to 2,300 °C (4,172 °F) |
| Oxyacetyene | Up to 3,300 °C (5,972 °F) |
Flame Temperature and Color
The color of a fire or a hot object offers a rough guide to its temperature. The glow emitted by a hot object is called black body radiation or incandescence. Its easily observed when heating a piece of metal:
- Deep red: 600-800 °C (1112-1800 °F)
- Orange-yellow: 1100 °C (2012 °F)
- White: 1300-1500 °C (2400-2700 °F)
- Blue: 1400-1650 °C (2600-3000 °F)
- Violet: 39400 °C (71000 °F)
However, flame color is not a good indicator of temperature because heated metal ions emit colored light. In other words, impurities in a fuel can produce a colored flame without increasing its temperature. For example, borax turns flames green, while potassium chloride turns fire violet.
Hottest Part of a Flame
The hottest part of a visible flame is blue, but science students are asked to use the top part of the flame for maximum heat. Why? The reason is that heat rises, so the top of the flame’s cone is a collection point for energy. Another reason to use the top of a flame is because its temperature is more consistent.
Hottest and Coolest Flames
Fire is always hot, but flames occur over a great temperature range. The coolest flame results from burning a regulated air-fuel mixture. This cool fire has flames with a temperature around 120° Celsius, which is still hotter than boiling water. The carbon subnitride (C4N4, also called diacetylene) and cyanogen-oxygen [(CN)2-O2] flames are hottest flames produced to date, with flame temperatures between 5000 and 6000 K (4727-5727 °C; 8540-10340 °F).. The extraordinary heat comes from from breaking the N2 triple bond and the high heats of formation of the compounds. While incredibly hot, these flame were blue-white and not violet.
References
- Jarosinski, Jozef; Veyssiere, Bernard (2009). Combustion Phenomena: Selected Mechanisms of Flame Formation, Propagation and Extinction. CRC Press. ISBN 0-8493-8408-7.
- Kirshenbaum, A. D.; A. V. Grosse (May 1956). “The Combustion of Carbon Subnitride, NC4N, and a Chemical Method for the Production of Continuous Temperatures in the Range of 5000–6000°K”. Journal of the American Chemical Society. 78 (9): 2020. doi:10.1021/ja01590a075
- Schmidt-Rohr, K (2015). “Why Combustions Are Always Exothermic, Yielding About 418 kJ per Mole of O2“. J. Chem. Educ. 92 (12): 2094–99. doi:10.1021/acs.jchemed.5b00333
- Thomas, N.; Gaydon, A. G.; Brewer, L. (1952). “Cyanogen Flames and the Dissociation Energy of N2“. The Journal of Chemical Physics. 20 (3): 369–374. doi:10.1063/1.1700426