Chemical Energy

Searching for Chemical Compounds

You have only to look around you, wherever you may be, to see chemical compounds by the thousands, everywhere; all are substances consisting of atoms of two or more elements held together by chemical bonds. Nearly every solid material on earth—nearly everything you see, animal, vegetable, or mineral—consists of chemical compounds, sometimes only one, sometimes several together. Solid substances composed of single elements uncombined with others are rarities; even "pure" gold and "pure" iron are almost never absolutely chemically pure.

When elements combine to form compounds, energy is either produced (or liberated—"evolved," as it is often called) or used up ("consumed"). A reaction that liberates energy, most often in the form of heat, is exothermic; a reaction that will take place only if energy is provided from some outside source is en-dothermic. Body warmth is produced by exothermic chemical reactions: without them, rigor mortis would soon set in. The earth's green plants, which directly or indirectly nourish almost all other living things, grow as a result of endothermic reactions that use sunlight as their energy source. Chemical reactions affect every moment of our lives: one could say that life is a series of chemical reactions, forever consuming and liberating energy.

Chemical reactions are also going on, all the time and everywhere, in the nonliving world around us. If you leave a steel chisel outside all winter, you know it will be rusty by spring. The iron in the steel has combined with atmospheric oxygen to form iron oxide (rust). The reaction is exothermic: about 5,100 joules of heat energy are liberated for every gram of rust formed,1 but the reaction happens too slowly for the heat to be noticeable except in carefully controlled laboratory experiments.

Chemical weathering of rock goes on all the time, as we saw in chapter 9. Minerals such as pyroxene, olivine, and hornblende, important ingredients of basalt, all contain iron that becomes oxidized—converted to rust—when fractured basalt is exposed to the air; rust is the chemical that makes iron ore red.

In the ordinary outdoor world, the commonest inorganic reactions in progress—those causing chemical weathering—attract much less attention than organic reactions taking place in living things. This is because weathering transforms a newly exposed rock surface so that it comes into equilibrium with its new environment, in contact with the atmosphere. The energy turnover in living material is much more obvious to human observers than are the very much slower inorganic reactions going on all the time around us. Both kinds of reactions consume and liberate energy: the contrast between them is in their speed.

The Energy in Chemical Bonds

Chemical reactions consist of the creation of new chemical compounds from existing ones. To create new compounds from old entails making and breaking chemical bonds; this is the point where energy enters the picture: energy is absorbed or released whenever chemical bonds change. Therefore, before continuing, we must consider what chemical bonds are and how they work.

Chemical bonds hold together the atoms in a molecule and, likewise, stick molecules to each other. A force must exist to create a bond, as the word "bond" implies. In chemical bonds, the force is electrical. It is a force of nature in the same way gravitation is a force of nature: both are fundamental characteristics of the physical world that cannot, at present, be explained in terms of anything more fundamental.

This does not mean, however, that chemical bonds are all alike. They differ from each other in magnitude, depending on which elements the bonded atoms belong to, and they differ in kind—bonds operate in several ways.

The simplest kind of bond, though not the commonest, is an ionic bond. The bonds holding together atoms of sodium and chlorine to make sodium chloride (table salt), for example, are of this kind: in a laboratory experiment, suppose a sodium atom that has become a positive ion by losing an electron comes close to a chlorine atom that has become a negative ion by picking up a stray electron; because of the electrical attraction between the two ions, one positively and the other negatively charged, they unite to form sodium chloride. Each step entails an energy change. The net result is the liberation of energy: for every gram of sodium chloride formed, 7,000 J of energy, in the form of heat, are produced.

The chemical bonds that hold together the atoms in organic molecules are known as covalent bonds. A covalent bond exists whenever two atoms share a pair of electrons. Covalent bonds are not confined to organic molecules, and the atoms bound together need not be of different elements. For example, in three common gases, hydrogen, oxygen, and nitrogen, each molecule consists of a pair of atoms of the element concerned, united by a covalent bond; that is why the gases are written, in chemical symbols, as H2, O2, and N2.

The simplest of all molecules, the hydrogen molecule, neatly illustrates the structure of a covalent bond. Each of its two atoms consists of one positively charged proton, the nucleus, and one negatively charged electron. The way they are covalently bound is diagrammed in figure 10.1. Figure 10.1a shows an isolated hydrogen atom; its single electron is somewhere in the circular cloud of dots surrounding the nucleus; one cannot say precisely where, just somewhere, moving rapidly, most frequently where the dots are densest. Figure 10.1b shows two hydrogen atoms united to form a molecule. The atoms share their two electrons: both are to be found somewhere in the oval cloud of dots encasing the two nuclei.

Free hydrogen in its natural state exists as molecules; to break all the bonds in a gram of hydrogen gas would require 217,000 J of energy. Breaking the bond in a single molecule uses 7.23 x 10-19 J. This is the covalent bond energy for one hydrogen molecule. Obviously joules are inconveniently big units for use with individual molecules; the appropriate unit for these tiny amounts is the electronvolt, abbreviated to eV (for more on electronvolts, see chapter 16). One eV is the same as 1.6 x 10-19 J; thus the bond energy for hydrogen is 4.5 eV. It is the energy required to break the bond in a hydrogen



Figure 10.1. (a) A solitary atom of hydrogen. (b) A molecule of hydrogen. The denser the stippling, the greater the chance of finding an electron there at any given time. The big black dots are the nuclei.

Figure 10.1. (a) A solitary atom of hydrogen. (b) A molecule of hydrogen. The denser the stippling, the greater the chance of finding an electron there at any given time. The big black dots are the nuclei.

molecule; equally, it is the energy liberated when two hydrogen atoms join to make a molecule.

Ionic bonds and covalent bonds normally hold together the atoms in a mol-ecule.2 Two other kinds of bonds attach molecules to each other; they are the bonds mentioned in chapter 9, which make solid objects stay solid, so that energy is needed to break them. They are not nearly as strong as the bonds joining the atoms in a molecule.

The intermolecular bonds that keep organic materials, including living organisms, from falling apart are hydrogen bonds; they form between a hydrogen atom in one molecule and an oxygen (or nitrogen or fluorine) atom in another. These three elements have atoms whose electrical characteristics enable them to make exceptionally strong bonds with hydrogen atoms. The commonest hydrogen bonds in living things are those in which oxygen is the element bonded to hydrogen.

Hydrogen bonds are the strongest intermolecular bonds, but they are not nearly as strong as covalent bonds, typically less than one-tenth as strong.3 Weaker bonds linking molecules to each other exist as well, known as van der Waals forces and caused by local concentrations of electric charge on molecular surfaces. They are much less common in natural materials than in synthetics such as plastics.4

Energy in Chemical Compounds

The energy liberated or consumed in any chemical reaction is the net result of the energies liberated and consumed when chemical bonds are made and broken. Complicated reactions like photosynthesis, for example, entail numerous makes and breaks to produce the end product, glucose, from its raw materials, water and carbon dioxide. The net result is consumption of energy; that is, the reaction is endothermic.

The energy consumed in photosynthesis is solar radiation, and the amount consumed is 16,000 J per gram of glucose created. The energy becomes stored in the glucose as chemical potential energy awaiting ultimate liberation as heat, in the same way that a rock poised at the top of a precipice stores gravitational potential energy awaiting liberation as movement. If you burn a gram of glucose, the photosynthetic reaction is reversed: water and carbon dioxide are produced and 16,000 J of energy are liberated—the potential energy is potential no longer.5 Or instead of burning it, you could eat it: the chemical reactions that constitute digestion would then yield 16,000 J of energy for the acts of living, such as moving, growing, and keeping warm.

We have already mentioned the energy liberated in two simple exothermic reactions; recall that just over 5,000 J are liberated for every gram of rust produced by the oxidation of iron; and 7,000 J are liberated when sodium and chlorine combine to form a gram of table salt. It would be necessary to supply the same amounts of energy from an outside source to undo these reactions: 5,000 J to unmake a gram of rust and 7,000 J to unmake a gram of table salt.

The energy change that accompanies every chemical reaction is called an enthalpy change. At first this term seems redundant: If you mean an energy change, why not say so? The reason is that an exothermic reaction produces two kinds of energy. The first is free energy, or energy capable of doing useful work—heating water, for instance. The second kind is entropy. Recall, from chapter 3, that it is never possible to make all the energy supplied to a system do useful work; some is always dissipated in a useless form, as entropy. We noted above that it takes 16,000 J of solar energy to energize the photosynthesis of one gram of glucose, and the same number of joules are liberated if the glucose is burned. But the liberated joules cannot all be useful energy, or we should have the makings of a perpetual motion machine. That is why the 16,000 J are given the name enthalpy, which is made up of both free energy (the active or useful kind) and entropy.6

Up to this point we have described entropy as "waste heat." Alternatively, it can be described as the energy of the random motion—the "milling about"— of the molecules in every substance whose temperature is greater than absolute zero (0 kelvins). The energy of this milling about depends on the chemical na-

ture of the substance as well as on its temperature. On average, the entropy of gases exceeds that of liquids, and the entropy of liquids exceeds that of solids. But this is true only on average. At room temperature, the entropy of solid table salt slightly exceeds that of liquid water; likewise the entropy of liquid alcohol exceeds that of helium gas. Solid substances have a wide range of entropies. For example, the entropy of lead is more than thirty times that of a diamond. This is because the atoms in a diamond are held firmly in their places in the diamond's crystal structure, whereas the atoms in a lump of lead are comparatively free to move; in brief, diamonds are well-knit, lead is rickety.

The change in enthalpy taking place in a chemical reaction is scarcely influenced by the temperature at which the reaction happens, but the partitioning of the enthalpy between free energy and entropy is strongly influenced: the ratio of free energy to entropy is much greater at low temperatures than at high ones.

Ice and Steam

Enthalpy changes are familiar to everybody. When water freezes, or vaporizes, the change in the water is a change in enthalpy even though no chemical reaction, in the usual sense, has taken place. Likewise when ice thaws or water vapor condenses.

Suppose you hold an ice cube in your hand: heat passes from your hand to the ice—that is why your hand is chilled—but the ice is not warmed above freezing point. On the contrary, it stays at freezing point while it gradually turns to water. The ice gains enthalpy, but no part of the enthalpy gained in this case is free energy; it is all entropy, manifested in the greatly increased mobility of the water molecules swirling around as a liquid instead of being held rigidly in place in crystalline ice.

Similarly, suppose you leave a pan of water boiling on a hot plate. Because the water is already boiling, its temperature will not rise any higher, but at the same time energy is passing from the hot plate to the water, changing the water's enthalpy. The change, an increase, has no free energy component; it consists wholly of an increase in entropy, manifested in the much greater mobility of the water molecules when they become a gas—the vapor rising from the boiling water—than they had as a liquid. If the water is boiled in a kettle with a loose lid, some of the enthalpy is free energy: it rattles the lid.

The change in enthalpy when a gram of ice melts is 335 J, and the change when a gram of water vaporizes is 2,259 J. Converting the joules to calories— the more familiar units in this context—gives 80 cal and 540 cal, respectively. These numbers will be recognized by many as the latent heat of freezing and the latent heat of vaporization of water. Indeed, the foregoing discussion deals with the same topic as the section on water vapor and energy transfers in chapter 5. Here we have seen how freezing and melting, or vaporizing and condensing, behave exactly like chemical reactions so far as energy transfers are concerned.

Chemical Energy to Electrical Energy and Back Again

To describe chemical reactions as endothermic or exothermic, according as they consume or liberate energy, is somewhat misleading: it suggests, falsely, that the energy involved when chemical bonds are made and broken always takes the form of heat. More inclusive terms are endergonic and exergonic. The former describes a reaction that must be supplied with free energy of some kind (not necessarily heat) from an external source if it is to proceed, the latter a reaction that proceeds without such a source. Note that it is the direction of flow of free energy, as opposed to enthalpy, that determines whether a reaction is endergonic or exergonic.7

Some chemical reactions entail the liberation or consumption of electrical energy. An obvious example is the charging and discharging of the lead storage battery in an automobile. This is the way it works.8 The battery consists of plates of pure metallic lead alternating with plates of lead oxide; there are spaces between the plates. All the lead plates are wired to a single conducting cable ending at the battery's negative terminal; likewise, all the lead oxide plates are wired to a cable ending at the positive terminal. The sheaf of spaced plates is immersed in a mixture of sulfuric acid and water.

Every time the circuit is completed (for example, when you switch on the ignition or the headlights), a chemical reaction starts in the battery. The lead plates react with the sulfuric acid to produce lead sulfate and free electrons, and the electrons start to stream through the circuit, producing an electric current. The flowing electrons do the required task—turning the starter motor, say, or lighting the headlights—and then return, via the positive terminal, to the lead oxide plates; there they combine with the lead oxide and some of the sulfuric acid to form more lead sulfate. When the battery is in use, lead sulfate accumulates on both kinds of plates while the sulfuric acid is gradually used up. As a result, the acid-and-water mixture becomes more dilute and, consequently, less dense; its density is what is measured when the battery is tested with a hydrometer.

When the acid becomes too dilute for the reaction to continue, the battery is said to be discharged. It can then be recharged by passing an electric current from some other source through it; this drives the chemical reactions in reverse, restoring things to their original condition: that is, the electrical energy supplied is converted back to chemical energy, available for reconversion to electrical energy when it is required.

The salient feature of what happens in an electrical storage battery is the transference of electrons from one material, lead, to another, lead oxide. Vast numbers of chemical reactions are of this kind: in simple ones, electrons are released by a pure element of one kind and become attached to a pure element of another kind. All these reactions—both simple and complicated—are called redox reactions.9 A redox reaction gives off energy when it goes in one direction (as when a battery is giving an electric current) and absorbs energy when it goes in the opposite direction (as when a battery is being recharged by an electric current fed into it). Redox reactions are what make living bodies live; biochemical redox reactions entail the transference of electrons between large, complicated organic molecules, but they are redox reactions nonetheless. They are what is happening when food is converted to energy; the electrical energy yielded by the reactions is transformed into the mechanical energy of movement and into heat.

All transfers of biochemical energy involve electron transfers, and they usually take place in several steps: the energy yielded by one reaction powers a second reaction that would not happen without it, which powers a third reaction, and so on. The chemical that functions most often as an intermediary in these sequences of reactions is adenosine triphosphate, well known by its acronym ATP; it has been called the principal carrier of biological energy.10

The electron transfers in biochemical reactions do not produce strong, easily detectable electric currents. Reactions having the same effects as the discharge and recharge of an automobile battery are much less common in nature than in the technological world. The only conspicuous examples in nature are the strong electric shocks delivered by certain species of fish when they are disturbed: the electric eel of South American rivers, the electric catfish of central Africa and the Nile, and several species of electric rays living in tropical and temperate seas in various parts of the world. All of them are capable of electrocuting a human being when their "batteries" are fully charged.

Getting Started With Solar

Getting Started With Solar

Do we really want the one thing that gives us its resources unconditionally to suffer even more than it is suffering now? Nature, is a part of our being from the earliest human days. We respect Nature and it gives us its bounty, but in the recent past greedy money hungry corporations have made us all so destructive, so wasteful.

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  • marja
    What force liberates energy in chemical reactions?
    7 years ago
  • Joona Rautavaara
    What is chemical energy in nature?
    7 years ago

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