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Chemical Energy

Long before the beginning of written history, humans knew how to start a fire and make use of the heat and light therefrom. Today we burn fuel for the production of chemical energy, which can drive electric generators to produce electrical energy that turns motors in factories and households.

It is a good time to think about the how, what, and why of chemical energy. First we ought to discuss what energy is. Energy is a measurement of the quantity of matter. Depending on which measuring instument is used, the quantity of matter in a sample will register as calories, watt hours, ergs, foot pounds, etc. The quantity of matter is also measured in units of mass, such as grams. The actual quantity of matter in any object is the count of its bits.

When matter is measured in terms of energy, it is only a partial measurement. Sometimes it is kinetic energy. It may be chemical potential energy or nuclear potential energy. However, the mass measurements are usually all-inclusive. Sometimes the measurement is partial, as in the additional mass due to acceleration.

It is somewhat misleading to say that a slice of bread has a hundred calories. What is really meant is that the entire process of bread being digested and reacting with water and oxygen and adenosine diphosphate will convert 100 calories of chemical potential energy into kinetic energy plus a miscellany of odds and ends. Similarly, in discussing fire, we are not examining the production of energy (which is nonsense), but the conversion of chemical potential energy into kinetic energy.

For a simple fuel we choose hydrogen. We do not measure the potential energy of a molecule of hydrogen. What we measure is the energy of the reaction between two molecules of hydrogen and one molecule of oxygen.

If we count bits in two hydrogen molecules plus one oxygen molecule, we have the starting energy of the system. The products of the reaction are two molecules of water. If we count the bits in two molecules of water, we discover that the sum is sligtly less for these two molecules than the sum for two hydrogen molecules plus one oxygen molecule. We can convert the number of missing bits into calories of potential chemical energy that has been converted into kinetic energy. (if we ignore the odds and ends) So where are the missing bits? They attached themselves to molecules of the surroundings.

In order to understand this, we must examine the mechanism of the reaction. Let the hydrogen gas enter the atmosphere through a narrow pipe. At room temperature the molecules of hydrogen, oxygen, and nitrogen intermingle without reacting.

Before we proceed, we would do well to know what temperature means. It is the average kinetic energy of the molecules. For one molecule, the average kinetic energy over time is the same as the average kinetic energy per molecule in a space, in an instant. At room temperatures , individual temperatures range from zero to a few hundred degrees. None of the molecules reaches kindling temperature.

Kindling temperature is the kinetic energy of one molecule at which it could strike one electron in some other molecule hard enough to cause it to escape from the molecule. As the electron escapes, it is accompanied by one of the nuclei. After all, the only thing that held the nucleus in the molecule was the attraction toward the electron.

When a flint rock is struck smartly with a piece of steel, tiny fragments fly off with a kinetic energy that corresponds to kindling temperature. These fragments strike air molecules hard enough to cause a few hdrogen atoms to fly free from their molecules. It turns out that hydrogen atoms are more strongly attracted to oxygen than oxygen atoms are attracted to each other. A free hydrogen atom flies toward an oxygen atom that is part of a pair. The speed of the hydrogen atom increases as it gets closer to its goal. In fact, the hydrogen atom gets too close and enters the area of repulsion. It then bounces off, loses speed,returns to the oxygen, bounces off, and repeats this dance forever. In the meantime, the second atom of the oxygen molecule loses its grip and departs. There follows a chain reaction and general pandemonium as all the hydrogen and oxygen atoms break loose and start dancing.

So far neither bits nor calories have departed. In fact the reaction is incomplete, because nothing is tied in place. (except nitrogen, which is a tough baby) However this will soon be remedied, bcause the active participants are surrounded by stable molecules of air. Each collision with an outsider transfers kinetic energy outward until the whole region is reduced to room temperature. When the newly formed water molecules settle down to inaction, the reaction is complete.

Then we can discover that there is a net loss of mass. But the missing matter is not lost. It is now part of the surroundings.

Now we ought to find out why the bond between hydrogen and oxygen is more attractive than the bond between hydrogen and hydrogen or between oxygen and oxygen.

The bond in hydrogen is due to the attraction of electron for proton. The two protons of the hydrogen molecule are less of an attraction than the eight protons of oxygen. Besides, one of the protons is itching to depart also. But the eight electrons in the oxygen should repel the hydrogen's electron. Not if it's escorted by a proton. Another factor is the ease with which eight electrons can rearrange themselves to make way for a newcomer.

In conclusion we find that no energy is produced by the burning of fuel. There is only the transfer of bits of matter from one arrangement to another. The benefit from the reaction comes from the use of heat and light by humans. It is not the quantity of heat that counts, but the temperature difference that makes it all worthwhile. But that is fuel for another lengthy essay.

Intermission


The author is Philip Mintz

E-mail: philmintz@lycos.com

Comments are welcome




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