Volume 1 — ENERGY

The Law of Conservation of Energy

We have now gone as far as we conveniently can in considering the intertwining strands of the atom and of electricity. It is time to turn to the third strand—energy.

To physicists the concept of “work” is that of exerting a force on a body and making it move through some distance. To lift a weight against the pull of gravity is work. To drive a nail into wood against the friction of its fibers is work.

Anything capable of performing work is said to possess “energy” from Greek words meaning “work within”. There are various forms of energy. Any moving mass possesses energy by virtue of its motion. That is, a moving hammer will drive a nail into wood, while the same hammer held motionlessly against the nailhead will not do so. Heat is a form of energy, since it will expand steam that will force wheels into motion that can then do work. Electricity, magnetism, sound, and light can be made to perform work and are forms of energy.

The forms of energy are so many and so various that scientists were eager to find some rule that covered them all and would therefore serve as a unifying bond. It did not seem impossible that such a rule might exist, since one had been found in connection with matter that appeared in even greater variety than energy did.

All matter, whatever its form and shape, possessed mass, and in the 1770s, the French chemist Antoine Laurent Lavoisier (1743-1794) discovered that the quantity of mass was constant. If a system of matter were isolated and made to undergo complicated chemical reactions, everything about it might change, but not its mass. A solid might turn into a gas; a single substance might change into two or three different substances, but whatever happened, the total mass at the end was exactly the same (as nearly as chemists could tell) as at the beginning. None was either created or destroyed, however, the nature of the matter might change. This was called the “law of conservation of mass”.

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Lavoisier in his laboratory during his studies on respiration. From a sketch made by Madame Lavoisier.

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Antoine Lavoisier and his wife.

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Naturally, it would occur to scientists to wonder if a similar law might hold for energy. The answer wasn’t easy to get. It wasn’t as simple to measure the quantity of energy as it was to measure the quantity of mass. Nor was it as simple to pen up a quantity of energy and keep it from escaping or from gaining additional quantity from outside, as it was in the case of mass.

Beginning in 1840, however, the English physicist James Prescott Joule (1818-1889) began a series of experiments in which he made use of every form of energy he could think of. In each case he turned it into heat and allowed the heat to raise the temperature of a given quantity of water. He used the rise in temperature as a measure of the energy. By 1847 he was convinced that any form of energy could be turned into fixed and predictable amounts of heat; that a certain amount of work was equivalent to a certain amount of heat.

In that same year, the German physicist Hermann Ludwig Ferdinand von Helmholtz (1821-1894) advanced the general notion that a fixed amount of energy in one form was equal to the same amount of energy in any other form. Energy might change its form over and over, but not change its amount. None could either be destroyed or created. This is the “law of conservation of energy”.

Chemical Energy

There is energy in a piece of wood. Left quietly to itself, it seems completely incapable of bringing about any kind of work. Set it on fire, however, and the wood plus the oxygen in the air will give off heat and light that are clearly forms of energy. The heat could help boil water and run a steam engine.

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The amount of energy in burning wood could be measured if it were mixed with air and allowed to burn in a closed container that was immersed in a known quantity of water. From the rise in temperature of the water, the quantity of energy produced could be measured in units called “calories” (from a Latin word for “heat”). The instrument was therefore called a “calorimeter”.

In the 1860s the French chemist Pierre Eugène Marcelin Berthelot (1827-1907) carried through hundreds of such determinations. His work and similar work by others made it clear that such “chemical energy”—the energy derived from chemical changes in matter—fit the law of conservation of energy.

Here’s how it looked in the last decades of the 19th century.

Molecules are composed of combinations of atoms. Within the molecules, the atoms stick together more or less tightly. It takes a certain amount of energy to pull a molecule apart into separate atoms against the resistance of the forces holding them together.

If, after being pulled apart, the atoms are allowed to come together again, they give off energy. The amount of energy they give off in coming together is exactly equal to the amount of energy they had to gain before they could separate.

This is true of all substances. For instance, hydrogen gas, as it is found on earth, is made up of molecules containing 2 hydrogen atoms each (H₂). Add a certain amount of energy and you pull the atoms apart; allow the atoms to come back together into paired molecules, and the added energy is given back again. The same is true for the oxygen molecule, which is made up of 2 oxygen atoms (O₂) and of the water molecule (H₂O). Always the amount of energy absorbed in one change is given off in the opposite change. The amount absorbed and the amount given off are always exactly equal.

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However, the amount of energy involved differs from molecule to molecule. It is quite hard to pull hydrogen molecules apart, and it is even harder to pull oxygen molecules apart. You have to supply about 12% more energy to pull an oxygen molecule apart than to pull a hydrogen molecule apart. Naturally, if you let 2 oxygen atoms come together to form an oxygen molecule, you get back 12% more energy than if you allow 2 hydrogen atoms to come together to form a hydrogen molecule.

It takes a considerably larger amount of energy to pull apart a water molecule into separate atoms than to pull apart either hydrogen or oxygen molecules. Naturally, that greater energy is also returned once the hydrogen and oxygen atoms are allowed to come back together into water molecules.

Next, imagine pulling apart hydrogen and oxygen molecules into hydrogen and oxygen atoms and then having those atoms come together to form water molecules. A certain amount of energy is put into the system to break up the hydrogen and oxygen molecules, but then a much greater amount of energy is given off when the water molecules form.

It is for that reason that a great deal of energy (mostly in the form of heat) is given off if a jet of hydrogen gas and a jet of oxygen gas are allowed to mix in such a way as to form water.

Just mixing the hydrogen and oxygen isn’t enough. The molecules of hydrogen and oxygen must be separated and that takes a little energy. The energy in a match flame is enough to raise the temperature of the mixture and to make the hydrogen and oxygen molecules move about more rapidly and more energetically. This increases the chance that some molecules will be broken up into separate atoms (though the actual process is rather complicated). An oxygen atom might then strike a hydrogen molecule to form water (O + H₂ → H₂O) and more energy is given off than was 53 absorbed from the match flame. The temperature goes up still higher so that further breakup among the oxygen and hydrogen molecules is encouraged.

The formation of a sodium chloride molecule.

This happens over and over again so that in very little time, the temperature is very high and the hydrogen and oxygen are combining to form water at an enormous rate. If a great deal of hydrogen and oxygen are well-mixed to begin with, the rate of reaction is so great that an explosion occurs.

Such a situation, in which each reacting bit of the system adds energy to the system by its reaction and brings about more reactions like itself, is called a “chain reaction”. Thus, a match flame put to one corner of a large sheet of paper will set that corner burning. The heat of the burning will ignite a 54 neighboring portion of the sheet and so on till the entire sheet is burned. For that matter a single smoldering cigarette end can serve to burn down an entire forest in a vastly destructive chain reaction.

Electrons and Energy

The discovery of the structure of the atom sharpened the understanding of chemical energy.

In 1904 the German chemist Richard Abegg (1869-1910) first suggested that atoms were held together through the transfer of electrons from one atom to another.

To see how this worked, one began by noting that electrons in an atom existed in a series of shells. The innermost shell could hold only 2 electrons, the next 8, the next 18 and so on. It turned out that some electron arrangements were more stable than others. If only the innermost shell contained electrons and it were filled with the 2 electrons that were all it could hold, then that was a stable arrangement. If an atom contained electrons in more than one shell and the outermost shell that held electrons held 8, that was a stable arrangement, too.

Thus, the helium atom has 2 electrons only, filling the innermost shell, and that is so stable an arrangement that helium undergoes no chemical reactions at all. The neon atom has 10 electrons—2 in the innermost shell, and 8 in the next—and it does not react. The argon atom has 18 electrons—2, 8, and 8—and it too is very stable.

But what if an atom did not have its electron shell so neatly filled. The sodium atom has 11 electrons—2, 8, and 1—while the fluorine atom has 9 electrons—2 and 7. If the sodium atom passed one of its electrons to a fluorine atom, both would have the stable configuration of neon—2 and 8. This, therefore, ought to have a great tendency to happen.

If it did happen, though, the sodium atom, minus 1 electron, would have a unit positive charge and would be Na⁺, 55 a positively charged ion. Fluorine with 1 electron in excess would become F⁻, a negatively charged ion. The 2 ions, with opposite charges, would cling together, since opposite charges attract, and thus the molecule of sodium fluoride (NaF) would be formed.

In 1916 the American chemist Gilbert Newton Lewis (1875-1946) carried this notion farther. Atoms could cling together not only as a result of the outright transfer of 1 or more electrons, but through sharing pairs of electrons. This sharing could only take place if the atoms remained close neighbors, and it would take energy to pull them apart and break up the shared pool, just as it would take energy to pull 2 ions apart against the attraction of opposite charges.

In this way the vague notions of atoms clinging together in molecules and being forced apart gave way to a much more precise picture of electrons being transferred or shared. The electron shifts could be dealt with mathematically by a system that came to be called “quantum mechanics” and chemistry was thus made a more exact science than it had ever been before.

The Energy of the Sun

The most serious problem raised by the law of conservation of energy involved the sun. Until 1847, scientists did not question sunlight. The sun radiated vast quantities of energy but that apparently was its nature and was no more to be puzzled over than the fact that the earth rotated on its axis.

Once Helmholtz had stated that energy could neither be created nor destroyed, however, he was bound to ask where the sun’s energy came from. It had, to man’s best knowledge, been radiating heat and light, with no perceptible change, throughout the history of civilization and, from what biologists and geologists could deduce, for countless ages earlier. Where, then, did that energy come from?

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The sun gave the appearance of being a huge globe of fire. Could it actually be that—a large heap of burning fuel, turning chemical energy into heat and light?

The sun’s mass was known and its rate of energy production was known. Suppose the sun’s mass were a mixture of hydrogen and oxygen and it were burning at a rate sufficient to produce the energy at the rate it was giving it off. If that were so, all the hydrogen and oxygen in its mass would be consumed in 1500 years. No chemical reaction in the sun could account for its having given us heat and light since the days of the pyramids, let alone since the days of the dinosaurs.

Was there some source of energy greater than chemical energy? What about the energy of motion? Helmholtz suggested that meteors might be falling into the sun at a steady rate. The energy of their collisions might then be converted into heat and light and this could keep the sun shining for as long as the supply of meteors held out—even millions of years.

This, however, would mean that the sun’s mass would be increasing steadily, and so would the force of its gravitational pull. With the sun’s gravitational field increasing steadily, the length of earth’s year would be decreasing at a measurable rate—but it wasn’t.

In 1854 Helmholtz came up with something better. He suggested that the sun was contracting. Its outermost layers were falling inward, and the energy of this fall was converted into heat and light. What’s more, this energy would be obtained without any change in the mass of the sun whatever.

Helmholtz calculated that the sun’s contraction over the 6000 years of recorded history would have reduced its diameter only 560 miles—a change that would not have been noticeable to the unaided eye. Since the development of the telescope, two and a half centuries earlier, the decrease in 57 diameter would have been only 23 miles and that was not measurable by the best techniques of Helmholtz’s day.

Working backward, however, it seemed that 25 million years ago, the sun must have been so large as to fill the earth’s orbit. Clearly the earth could not then have existed. In that case, the maximum age of the earth was only 25 million years.

Geologists and biologists found themselves disturbed by this. The slow changes in the earth’s crust and in the evolution of life made it seem very likely that the earth must have been in existence—with the sun delivering heat and light very much in the present fashion—for many hundreds of millions of years.

Yet there seemed absolutely no other way of accounting for the sun’s energy supply. Either the law of conservation of energy was wrong (which seemed unlikely), or the painfully collected evidence of geologists and biologists was wrong (which seemed unlikely),—or there was some source of energy greater than any known in the 19th century, whose existence had somehow escaped mankind (which also seemed unlikely).

Yet one of those unlikely alternatives would have to be true. And then in 1896 came the discovery of radioactivity.

The Energy of Radioactivity

It eventually became clear that radioactivity involved the giving off of energy. Uranium emitted gamma rays that we now know to be a hundred thousand times as energetic as ordinary light rays. What’s more, alpha particles were being emitted at velocities of perhaps 30,000 kilometers per second, while the lighter beta particles were being shot off at velocities of up to 250,000 kilometers per second (about 0.8 times the velocity of light).

At first, the total energy given off by radioactive substances seemed so small that there was no use worrying 58 about it. The amount of energy liberated by a gram of uranium in 1 second of radioactivity was an insignificant fraction of the energy released by a burning candle.

In a few years, however, something became apparent. A lump of uranium might give off very little energy in a second, but it kept on for second after second, day after day, month after month, and year after year with no perceptible decrease. The energy released by the uranium over a very long time grew to be enormous. It eventually turned out that while the rate at which uranium delivered energy did decline, it did so with such unbelievable slowness that it took 4.5 billion years (!) for that rate to decrease to half what it was to begin with.

If all the energy delivered by a gram of uranium in the course of its radioactivity over many billions of years was totalled, it was enormously greater than the energy produced by the burning of a candle with a mass equal to that of uranium.

Let’s put it another way. We might think of a single uranium atom breaking down and shooting off an alpha particle. We might also think of a single carbon atom combining with 2 oxygen atoms to form carbon dioxide. The uranium atom would give off 2,000,000 times as much energy in breaking down, as the carbon atom would in combining.

The energy of radioactivity is millions of times as intense as the energy released by chemical reactions. The reason mankind had remained unaware of radioactivity and very aware of chemical reactions was, first, that the most common radioactive processes are so slow that their great energies were stretched over such enormous blocks of time as to be insignificant on a per second basis.

Secondly, chemical reactions are easily controlled by changing quantities, concentrations, temperatures, pressures, states of mixtures, and so on, and this makes them easy to 59 take note of and to study. The rate of radioactive changes, however, could not apparently be altered. The early investigators quickly found that the breakdown of uranium-238, for instance, could not be hastened by heat, pressure, changes in chemical combination, or, indeed, anything else they could think of. It remained incredibly slow.

But despite all this, radioactivity had at last been discovered and the intensity of its energies was recognized and pointed out in 1902 by Marie Curie and her husband Pierre Curie (1859-1906).

Where, then, did the energy come from? Could it come from the outside? Could the radioactive atoms somehow collect energy from their surroundings, concentrate it several million-fold, and then let it out all at once?

To concentrate energy in this fashion would violate something called “the second law of thermodynamics”. This was first proposed in 1850 by the German physicist Rudolf Julius Emmanuel Clausius (1822-1888) and had proved so useful that physicists did not like to abandon it unless they absolutely had to.

Another possibility was that radioactive atoms were creating energy out of nothing. This, of course, violated the law of conservation of energy (also called “the first law of thermodynamics”) and physicists preferred not to do that either.

The only thing that seemed to remain was to suppose that somewhere within the atom was a source of energy that had never made itself evident to humanity until the discovery of radioactivity. Becquerel was one of the first to suggest this.

It might have seemed at first that only radioactive elements had this supply of energy somewhere within the atom, but in 1903 Rutherford suggested that all atoms had a vast energy supply hidden within themselves. The supply in uranium and thorium leaked slightly, so to speak, and that was all that made them different.

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The room in which the Curies discovered radium. Pierre Curie’s writing is on the blackboard.

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But if a vast supply of energy existed in atoms, it was possible that the solution to the puzzle of the sun’s energy might rest there. As early as 1899 the American geologist Thomas Chrowder Chamberlin (1843-1928) was already speculating about a possible connection between radioactivity and the sun’s energy.

If it were some variety of this newly discovered source of energy (not necessarily ordinary radioactivity, of course) that powered the sun—millions of times as intense as chemical energy—then the sun might be pouring out energy for hundreds of millions of years without perceptible physical change—just as uranium would show scarcely any change even in so mighty a time span. The sun would not have to be contracting; it would not have had to fill the earth’s orbit 25,000,000 years ago.

This was all exciting, but in 1900 the structure of the atom had not yet been worked out and this new energy was just a vague supposition. No one had any idea of what it actually might be or where in the atom it might be located. It could only be spoken of as existing “within the atom” and was therefore called “atomic energy”. Through long habit, it is still called that much of the time. And yet “atomic energy” is not a good name. In the first couple of decades of the 20th century, it became apparent that ordinary chemical energy involved electron shifts and those electrons were certainly components of atoms. This meant that a wood fire was a kind of atomic energy.

The electrons, however, existed only in the outer regions of the atom. Once Rutherford worked out the theory of the nuclear atom, it became apparent that the energy involved in radioactivity and in solar radiation had to involve components of the atom that were more massive and more energetic than the light electrons. The energy had to come, somehow, from the atomic nucleus.

What is involved then in radioactivity and in the sun is “nuclear energy”. That is the proper name for it and in the 62 next section we will consider the subsequent history of the nuclear energy that broke upon the startled consciousness of scientists as the 20th century opened and which, less than half a century later, was to face mankind with untold consequences for good and for evil.