# How to thermodynamically understand process of burning a piece of coal?

Let's imagine that I have a match in hand and nugget of coal on my desk. Then I light up the match and place it for few seconds near the coal so a tiny piece of nugget catches fire. Then another piece catches fire, then another and soon all the nugget is burnt down.

How did it happen? I gave the nugget just enough heat to burn the first piece. Where does come energy to burn the rest of nugget from?

• When you light a nugget of coal, how does the flame persist even after you've extinguished the match? Commented Feb 17, 2015 at 18:12
• @lemon Because I gave it some external energy which was equal to the enthalpy of combustion? But it was only equal to the enthalpy of a little piece of the nugget. I'm confused. Commented Feb 17, 2015 at 18:20
• @lemon pardon my stupid question, but my thermodynamics course was based mainly on memorizing very long derivations. Worst part of my life. Commented Feb 17, 2015 at 18:38
• The nuggets release heat + flammable gases when they burn, which allows other nuggets to catch fire. Commented Feb 17, 2015 at 19:11
• @pentane But why do I need to deliver heat before the nugget releases more heat? I guess it is because of energy of bondings, but I would like to see more mathematical answer. Commented Feb 17, 2015 at 19:19

## 2 Answers

It is called combustion, and it happens in materials which have a lower energy content when their component molecules join with the oxygen in the atmosphere, than when in a solid/liquid structure. When energy is given to start the fire the piece of coal burns and releases energy with excess enough to sustain the reaction and leave heat energy for use.

Combustion is a high-temperature exothermic chemical reaction between a fuel and an oxidant, usually atmospheric oxygen, that produces oxidized, often gaseous products, in a mixture termed as smoke.

• But why do I need to give energy to start the fire, while even without my intervention the gaseous products have lower energy? Why isn't it spontaneous? And could it be understood only with classical thermodynamics, without molecular interpretation of matter? Commented Feb 17, 2015 at 19:29
• Oh! I found it in a thermodynamics textbook! Need of initial flame is not the matter of thermodynamics, but of reaction kinetics! But they didn't explained why and how. Commented Feb 17, 2015 at 19:34
• The joining of oxygen to carbon is not spontaneous at room temperatures, but at temperatures high enough to start breaking the chemical bonds, become gaseous to mix with oxygen. en.wikipedia.org/wiki/Spontaneous_combustion en.wikipedia.org/wiki/Fire_point Commented Feb 17, 2015 at 19:34
• Now, with your help I can better understand what's in my handbook. Thanks! Commented Feb 17, 2015 at 19:38

Burning/combustion involves chemical reactions that transform the molecules of the initial compounds into molecules with lower (i.e., bigger) binding energy. In the case of burning coal we are mainly speaking about breaking the crabon bonds of the hydrocarbons in coal (see here for an example) and those in the atmospheric oxygen, $$O_2$$, whereas the final compounds are mainly $$CO$$ and $$CO_2$$.

The difference in the binding energies is the heat released, and can be quite significant. However, the initial carbon and oxygen bonds are already rather strong and do not spontaneously break at room temperature. Therefore one needs to heat the carbon and the oxygen, to break the bonds and initiate the reaction. Once the reaction is launched, the heat released through it breaks additional bonds and thus the reaction continues.

Remarks:

• One could point out additional layers of complexity: e.g., it is not the carbon itself that burns, but the gases released from it, which also requires heat.
• If the description above does not immediately looks like a thermodynamic one, it is because here we are not talking about a reaction running in equilibrium conditions, but rather about a dynamic transformation. On otehr words, we are dealing with non-equilibrium thermodynamics / physical kinetics.