CHEMICAL REACTIONS

CONCEPT

If chemistry were compared to a sport, then the study of atomic and molecular properties, along with learning about the elements and how they relate on the periodic table, would be like going to practice. Learning about chemical reactions, which includes observing them and sometimes producing them in a laboratory situation, is like stepping out onto the field for the game itself. Just as every sport has its "vocabulary"—the concepts of offense and defense, as well as various rules and strategies—the study of chemical reactions involves a large set of terms. Some aspects of reactions may seem rather abstract, but the effects are not. Every day, we witness evidence of chemical reactions—for instance, when a fire burns, or metal rusts. To an even greater extent, we are surrounded by the products of chemical reactions: the colors in the clothes we wear, or artificial materials such as polymers, used in everything from nylon running jackets to plastic milk containers.

HOW IT WORKS

What Is a Chemical Reaction?

If liquid water is boiled, it is still water; likewise frozen water, or ice, is still water. Melting, boiling, or freezing simply by the application of a change in temperature are examples of physical changes, because they do not affect the internal composition of the item or items involved. A chemical change, on the other hand, occurs when the actual composition changes—that is, when one substance is transformed into another. Water can be chemically changed, for instance, when an electric current is run through a sample, separating it into oxygen and hydrogen gas.

Chemical change requires a chemical reaction, a process whereby the chemical properties of a substance are altered by a rearrangement of the atoms in the substance. Of course we cannot see atoms with the naked eye, but fortunately, there are a number of clues that tell us when a chemical reaction has occurred. In many chemical reactions, for instance, the substance may experience a change of state or phase—as for instance when liquid water turns into gaseous oxygen and hydrogen as a result of electrolysis.

HOW DO WE KNOW WHEN A CHEMICAL REACTION HAS OCCURRED?

Changes of state may of course be merely physical—as for example when liquid water is boiled to form a vapor. (These and other examples of physical changes resulting from temperature changes are discussed in the essays on Properties of Matter; Temperature and Heat.) The vapor produced by boiling water, as noted above, is still water; on the other hand, when liquid water is turned into the elemental gases hydrogen and oxygen, a more profound change has occurred.

Likewise the addition of liquid potassium chromate (K₂CrO₄) to a solution of barium nitrate (Ba[NO₃]₂ forms solid barium chromate (BaCrO₄). In the reaction described, a solution is also formed, but the fact remains that the mixture of two solids has resulted in the formation of a solid in a different solution. Again, this is a far more complex phenomenon than the mere freezing of water to form ice: here the fundamental properties of the materials involved have changed.

The physical change of water to ice or steam, of course, involves changes in temperature; likewise, chemical changes are often accompanied by changes in temperature, the crucial difference being that these changes are the result of alterations in the chemical properties of the substances involved. Such is the case, for instance, when wood burns in the presence of oxygen: once wood is turned to ash, it has become an entirely different mixture than it was before. Obviously, the ashes cannot be simply frozen to turn them back into wood again. This is an example of an irreversible chemical reaction.

Chemical reactions may also involve changes in color. In specific proportions and under the right conditions, carbon—which is black—can be combined with colorless hydrogen and oxygen to produce white sugar. This suggests another kind of change: a change in taste. (Of course, not every product of a chemical reaction should be tasted—some of the compounds produced may be toxic, or at the very least, extremely unpleasant to the taste buds.) Smell, too, can change. Sulfur is odorless in its elemental form, but when combined with hydrogen to form hydrogen sulfide (H₂S), it becomes an evil-smelling, highly toxic gas.

The bubbling of a substance is yet another clue that a chemical reaction has occurred. Though water bubbles when it boils, this is merely because heat has been added to the water, increasing the kinetic energy of its molecules. But when hydrogen peroxide bubbles when exposed to oxygen, no heat has been added. As with many of the characteristics of a chemical reaction described above, bubbling does not always occur when two chemicals react; however, when one of these clues is present, it tells us that a chemical reaction may have taken place.

REAL-LIFE APPLICATIONS

Chemical Equations

In every chemical reaction, there are participants known as reactants, which, by chemically reacting to one another, result in the creation of a product or products. As stated earlier, a chemical reaction involves changes in the arrangement of atoms. The atoms in the reactants (or, if the reactant is a compound, the atoms in its molecules) are rearranged. The atomic or molecular structure of the product is different from that of either reactant.

Note, however, that the number of atoms does not change. Atoms themselves are neither created nor destroyed, and in a chemical reaction, they merely change partners, or lose partners altogether as they return to their elemental form. This is a critical principle in chemistry, one that proves that medieval alchemists' dream of turning lead into gold was based on a fallacy. Lead and gold are both elements, meaning that each has different atoms. To imagine a chemical reaction in which one becomes the other is like saying "one plus one equals one."

SYMBOLS IN A CHEMICAL EQUATION.

In a mathematical equation, the sums of the numbers on one side of the equals sign must be the same as the sum of the numbers on the other side. The same is true of a chemical equation, a representation of a chemical reaction in which the chemical symbols on the left stand for the reactants, and those on the right are the product or products. Instead of an equals sign separating them, an arrow, pointing to the right to indicate the direction of the reaction, is used.

Chemical equations usually include notation indicating the state or phase of matter for the reactants and products. These symbols are as follows:

• (s) : solid
• (l) : liquid
• (g) : gas
• (aq) : dissolved in water (an aqueous solution)

The fourth symbol, of course, does not indicate a phase of matter per se (though obviously it appears to be a liquid); but as we shall see, aqueous solutions play a role in so many chemical reactions that these have their own symbol. At any rate, using this notation, we begin to symbolize the reaction of hydrogen and oxygen to form water thus: H(g) + O(g) →H₂O(l).

This equation as written, however, needs to be modified in several ways. First of all, neither hydrogen nor oxygen is monatomic. In other words, in their elemental form, neither appears as a single atom; rather, these form diatomic (two-atom) molecules. Therefore, the equation must be rewritten as H₂(g) + O₂(g) →H₂O(l). But this is still not correct, as a little rudimentary analysis will show.

Balancing Chemical Equations

When checking a chemical equation, one should always break it down into its constituent elements, to determine whether all the atoms on the left side reappear on the right side; otherwise, the result may be an incorrect equation, along the lines of "1 + 1 = 1." That is exactly what has happened here. On the left side, we have two hydrogen atoms and two oxygen atoms; on the right side, however, there is only one oxygen atom to go with the two hydrogens.

Obviously, this equation needs to be corrected to account for the second oxygen atom, and the best way to do that is to show a second water molecule on the right side. This will be represented by a 2 before the H₂O, indicating that two water molecules now have been created. The 2, or any other number used for showing more than one of a particular chemical species in a chemical equation, is called a coefficient. Now we have H₂(g) + O₂(g) →2H₂O(l).

Is this right? Once again, it is time to analyze the equation, to see if the number of atoms on the left equals the number on the right. Such analysis can be done in a number of ways: for instance, by symbolizing each chemical species as a circle with chemical symbols for each element in it. Thus a single water molecule would be shown as a circle containing two H's and one O.

Whatever the method used, analysis will reveal that the problem of the oxygen imbalance has been solved: now there are two oxygens on the left, and two on the right. But solving that problem has created another, because now there are four hydrogen atoms on the right, as compared with two on the left. Obviously, another coefficient of 2 is needed, this time in front of the hydrogen molecule on the left. The changed equation is thus written as: 2H₂(g) + O₂(g) → 2H₂O(l). Now, finally, the equation is correct.

THE PROCESS OF BALANCING CHEMICAL EQUATIONS.

What we have done is to balance an unbalanced equation. An unbalanced equation is one in which the numbers of atoms on the left are not the same as the number of atoms on the right. Though an unbalanced equation is incorrect, it is sometimes a necessary step in the process of finding the balanced equation—one in which the number of atoms in the reactants and those in the product are equal.

In writing and balancing a chemical equation, the first step is to ascertain the identities, by formula, of the chemical species involved, as well as their states of matter. After identifying the reactants and product, the next step is to write an unbalanced equation. After that, the unbalanced equation should be subjected to analysis, as demonstrated above.

The example used, of course, involves a fairly simple substance, but often, much more complex molecules will be part of the equation. In performing analysis to balance the equation, it is best to start with the most complex molecule, and determine whether the same numbers and proportions of elements appear in the product or products. After the most complicated molecule has been dealt with, the second-most complex can then be addressed, and so on.

Assuming the numbers of atoms in the reactant and product do not match, it will be necessary to place coefficients before one or more chemical species. After this has been done, the equation should again be checked, because as we have seen, the use of a coefficient to straighten out one discrepancy may create another. Note that only coefficients can be changed; the formulas of the species themselves (assuming they were correct to begin with) should not be changed.

After the equation has been fully balanced, one final step is necessary. The coefficients must be checked to ensure that the smallest integers possible have been used. Suppose, in the above exercise, we had ended up with an equation that looked like this: 12H₂(g) + 6O₂(g) →12H₂O(l). This is correct, but not very "clean." Just as a fraction such as 12/24 needs to be reduced to its simplest form, 1/2, the same is true of a chemical equation. The coefficients should thus always be the smallest number that can be used to yield a correct result.

Types of Chemical Reactions

Note that in chemical equations, one of the symbols used is (aq), which indicates a chemical species that has been dissolved in water—that is, an aqueous solution. The fact that this has its own special symbol indicates that aqueous solutions are an important part of chemistry. Examples of reactions in aqueous solutions are discussed, for instance, in the essays on Acid-Base Reactions; Chemical Equilibrium; Solutions.

Another extremely important type of reaction is an oxidation-reduction reaction. Sometimes called a redox reaction, an oxidation-reduction reaction occurs during the transfer of electrons. The rusting of iron is an example of an oxidation-reduction reaction; so too is combustion. Indeed, combustion reactions—in which oxygen produces energy so rapidly that a flame or even an explosion results—are an important subset of oxidation-reduction reactions.

REACTIONS THAT FORM WATER, SOLIDS, OR GASES.

Another type of reaction is an acid-base reaction, in which an acid is mixed with a base, resulting in the formation of water along with a salt.

Other reactions form gases, as for instance when water is separated into hydrogen and oxygen. Similarly, heating calcium carbonate (lime-stone) to make calcium oxide or lime for cement also yields gaseous carbon dioxide: CaCO³(s) + heat →CaO(s) + CO²(g).

There are also reactions that form a solid, such as the one mentioned much earlier, in which solid BaCrO₄(s) is formed. Such reactions are called precipitation reactions. But this is also a reaction in an aqueous solution, and there is another product: 2KNO₃(aq), or potassium nitrate dissolved in water.

SINGLE AND DOUBLE DISPLACEMENT.

The reaction referred to in the preceding paragraph also happens to be an example of another type of reaction, because two anions (negatively charged ions) have been exchanged. Initially K⁺ and CrO₄²⁻ were together, and these reacted with a compound in which Ba²⁺ and NO₃⁻ were combined. The anions changed places, an instance of a double-displacement reaction, which is symbolized thus: AB + CD →AD + CB.

It is also possible to have a single-displacement reaction, in which an element reacts with a compound, and one of the elements in the compound is released as a free element. This can be represented symbolically as A + BC →B + AC. Single-displacement reactions often occur with metals and with halogens. For instance, a metal(A) reacts with an acid (BC) to produce hydrogen (B) and a salt (AC).

COMBINATION AND DECOMPOSITION.

A synthesis, or combination, reaction is one in which a compound is formed from simpler materials—whether those materials be elements or simple compounds. A basic example of this is the reaction described earlier in relation to chemical equations, when hydrogen and oxygen combine to form water. On the other hand, some extremely complex substances, such as the polymers in plastics and synthetic fabrics such as nylon, also involve synthesis reactions.

When iron rusts (in other words, it oxidizes in the presence of air), this is both an oxidation-reduction and a synthesis reaction. This also represents one of many instances in which the language of science is quite different from everyday language. If a piece of iron—say, a railing on a balcony—rusts due to the fact that the paint has peeled off, it would seem from an unscientific standpoint that the iron has "decomposed." However, rust (or rather, metal oxide) is a more complex substance than the iron, so this is actually a synthesis or combination reaction.

A true decomposition reaction occurs when a compound is broken down into simpler compounds, or even into elements. When water is subjected to electrolysis such that the hydrogen and oxygen are separated, this is a decomposition reaction. The fermentation of grapes to make wine is also a form of decomposition.

And then, of course, there are the processes that normally come to mind when we think of "decomposition": the decay or rotting of a formerly living thing. This could also include the decay of something, such as an item of food, made from a formerly living thing. In such instances, an organic substance is eventually broken down through a number of processes, most notably the activity of bacteria, until it ultimately becomes carbon, nitrogen, oxygen, and other elements that are returned to the environment.

SOME OTHER PARAMETERS.

Obviously, there are numerous ways to classify chemical reactions. Just to complicate things a little more, they can also be identified as to whether they produce heat (exothermic) or absorb heat (endothermic). Combustion is clearly an example of an exothermic reaction, while an endothermic reaction can be exemplified by the process that takes place in a cold pack. Used for instance to prevent swelling on an injured ankle, a cold pack contains an ampule that absorbs heat when broken.

Still another way to identify chemical reactions is in terms of the phases of matter involved. We have already seen that some reactions form gases, some solids, and some yield water as one of the products. If reactants in one phase of matter produce a substance or substances in the same phase (liquid, solid, or gas), this is called a homogeneous reaction. On the other hand, if the reactants are in different phases of matter, or if they produce a substance or substances that are in a different phase, this is called a heterogeneous reaction.

An example of a homogeneous reaction occurs when gaseous nitrogen combines with oxygen, also a gas, to produce nitrous oxide, or "laughing gas." Similarly, nitrogen and hydrogen combine to form ammonia, also a gas. But when hydrogen and oxygen form water, this is a heterogeneous reaction. Likewise, when a metal undergoes an oxidation-reduction reaction, a gas and a solid react, resulting in a changed form of the metal, along with the production of new gases.

Finally, a chemical reaction can be either reversible or irreversible. Much earlier, we described how wood experiences combustion, resulting in the production of ash. This is clearly an example of an irreversible reaction. The atoms in the wood and the air that oxidized it have not been destroyed, but it would be impossible to put the ash back together to make a piece of wood. By contrast, the formation of water by hydrogen and oxygen is reversible by means of electrolysis.

KEEPING IT ALL STRAIGHT.

The different classifications of reactions discussed above are clearly not mutually exclusive; they simply identify specific aspects of the same thing. This is rather like the many physical characteristics that describe a person: gender, height, weight, eye color, hair color, race, and so on. Just because someone is blonde, for instance, does not mean that the person cannot also be brown-eyed; these are two different parameters that are more or less independent.

On the other hand, there is some relation between these parameters in specific instances: for example, females over six feet tall are rare, simply because women tend to be shorter than men. But there are women who are six feet tall, or even considerably taller. In the same way, it is unlikely that a reaction in an aqueous solution will be a combustion reaction—yet it does happen, as for instance when potassium reacts with water.

Studying Chemical Reactions

Several aspects or subdisciplines of chemistry are brought to bear in the study of chemical reactions. One is stoichiometry (stoy-kee-AH-muh-tree), which is concerned with the relationships among the amounts of reactants and products in a chemical reaction. The balancing of the chemical equation for water earlier in this essay is an example of basic stoichiometry.

Chemical thermodynamics is the area of chemistry that addresses the amounts of heat and other forms of energy associated with chemical reactions. Thermodynamics is also a branch of physics, but in that realm, it is concerned purely with physical processes involving heat and energy. Likewise physicists study kinetics, associated with the movement of objects. Chemical kinetics, on the other hand, involves the study of the collisions between molecules that produce a chemical reaction, and is specifically concerned with the rates and mechanisms of reaction.

SPEEDING UP A CHEMICAL REACTION.

Essentially, a chemical reaction is the result of collisions between molecules. According to this collision model, if the collision is strong enough, it can break the chemical bonds in the reactants, resulting in a rearrangement of the atoms to form products. The more the molecules collide, the faster the reaction. Increase in the numbers of collisions can be produced in two ways: either the concentrations of the reactants are increased, or the temperature is increased. In either case, more molecules are colliding.

Increases of concentration and temperature can be applied together to produce an even faster reaction, but rates of reaction can also be increased by use of a catalyst, a substance that speeds up the reaction without participating in it either as a reactant or product. Catalysts are thus not consumed in the reaction. One very important example of a catalyst is an enzyme, which speeds up complex reactions in the human body. At ordinary body temperatures, these reactions are too slow, but the enzyme hastens them along. Thus human life can be said to depend on chemical reactions aided by a wondrous form of catalyst.

WHERE TO LEARN MORE

Bender, Hal. "Chemical Reactions." Clackamas Community College (Web site). <http://dl.clackamas.cc.or.us/ch104-01/chemical.htm> (June 3, 2001).

"Catalysis, Separations, and Reactions." Accelrys (Web site). <http://www.accelrys.com/chemicals/catalysis/> (June 3, 2001).

Goo, Edward. "Chemical Reactions" (Web site). <http://www-classes.usc.edu/engr/ms/125/MDA125/reactions/> (June 3, 2001).

Knapp, Brian J. Oxidation and Reduction. Illustrated by David Woodroffe. Danbury, CT: Grolier Educational, 1998.

Knapp, Brian J. Energy and Chemical Change. Illustrated by David Woodroffe. Danbury, CT: Grolier Educational, 1998.

Newmark, Ann. Chemistry. New York: Dorling Kindersley, 1993.

"Periodic Table: Chemical Reaction Data." WebElements (Web site). <http://www.webelements.com/webelements/elements/text/periodic-table/chem.html>(June 3, 2001).

Richards, Jon. Chemicals and Reactions. Brookfield, CT: Copper Beech Books, 2000.

"Types of Chemical Reactions" (Web site). <http://www.usoe.k12.ut.us/curr/science/sciber00/8th/matter/sciber/chemtype.htm> (June 3, 2001).

Zumdahl, Steven S. Introductory Chemistry: A Foundation, 4th ed. Boston: Houghton Mifflin, 2000.

KEY TERMS

ACID-BASE REACTION:

A chemical reaction in which an acid is mixed with a base, resulting in the formation of water along with a salt.

AQUEOUS SOLUTIONS:

A mixture of water and any substance that is solvent in it.

BALANCED EQUATION:

A chemical equation in which the numbers of atoms in the reactants and those in the product areequal. In the course of balancing an equation, coefficients may need to be applied to one or more of the chemical species involved; however, the actual formulas of the species cannot be changed.

CATALYST:

A substance that speeds upa chemical reaction without participating in it either as a reactant or product. Catalysts are thus not consumed in the reaction.

CHEMICAL EQUATION:

A representation of a chemical reaction in which the chemical symbols on the left stand for the reactants, and those on the right for the product or products. On paper, a chemical equation looks much like a mathematical one; however, instead of an equals sign, a chemical equation uses an arrow to show the direction of the reaction.

CHEMICAL KINETICS:

the study of the rate at which chemical reactions occur.

CHEMICAL REACTION:

A process whereby the chemical properties of a substance are changed by a rearrangement of the atoms in the substance.

CHEMICAL SPECIES:

A generic term used for any substance studied in chemistry—whether it be an element, compound, mixture, atom, molecule, ion, and so forth.

CHEMICAL THERMODYNAMICS:

The study of the amounts of heat and other forms of energy associated with chemical reactions.

COEFFICIENT:

A number used to indicate the presence of more than one unit—typically, more than one molecule—of a chemical species in a chemical equation. For instance, 2H₂O indicates two water molecules. (Note that 1 is never used as a coefficient.)

COLLISION MODEL:

The theory that chemical reactions are the result of collisions between molecules that are strong enough to break bonds in the reactants, resulting in a rearrangement of atoms to form a product or products.

DECOMPOSITION REACTION:

A chemical reaction in which a compound is broken down into simpler compounds, or even into elements. This is the opposite of a synthesis or combination reaction.

DOUBLE-DISPLACEMENT REACTION:

A chemical reaction in which the partners in two compounds changeplaces. This can be symbolized as AB + CD →AD + CB. Compare single-displacement reaction.

ENDOTHERMIC:

A term describing a chemical reaction in which heat is absorbed or consumed.

EXOTHERMIC:

A term describing a chemical reaction in which heat is produced.

HETEROGENEOUS:

A term describing a chemical reaction in which the reactants are in different phases of matter (liquid, solid, or gas), or one in which the product is in a different phase from that of the reactants.

HOMOGENEOUS:

A term describing a chemical reaction in which the reactants and the product are all in the same phase of matter (liquid, solid, or gas).

OXIDATION-REDUCTION REACTION:

A chemical reaction involving the transfer of electrons.

PRECIPITATION REACTION:

A chemical reaction in which a solid isformed.

PRODUCT:

The substance or substances that result from a chemical reaction.

REACTANT:

A substance that interacts with another substance in a chemical reaction, resulting in the formation of aproduct.

SINGLE-DISPLACEMENT REACTION:

A chemical reaction in which an element reacts with a compound, and one of the elements in the compound is released as a free element. This can be represented symbolically as A + BC →B + AC. Compare double-displacement reaction.

STOICHIOMETRY:

The study of the relationships among the amounts of reactants and products in a chemical reaction. Producing a balanced equation requires application of stoichiometry (pronounced "stoy-kee-AH-muh-tree").

SYNTHESIS OR COMBINATIONREACTION:

A chemical reaction in which a compound is formed from simpler materials—either elements or simple compounds. It is the opposite of a decomposition reaction.

UNBALANCED EQUATION:

A chemical equation in which the sum of atoms in the product or products does not equal the sum of atoms in the reactants. Initial observations of a chemical reaction usually produce an unbalanced equation, which needs to be analyzed and corrected (by the use of coefficients) to yield a balancedequation.

Chemical Reactions

A chemical reaction is a process in which one set of chemical substances (reactants) is converted into another (products). It involves making and breaking chemical bonds and the rearrangement of atoms. Chemical reactions are represented by balanced chemical equations, with chemical formulas symbolizing reactants and products. For specific chemical reactants, two questions may be posed about a possible chemical reaction. First, will a reaction occur? Second, what are the possible products if a reaction occurs? This

entry will focus only on the second question. The most reliable answer is obtained by conducting an experiment—mixing the reactants and then isolating and identifying the products. We can also use periodicity, since elements within the same group in the Periodic Table undergo similar reactions. Finally, we can use rules to help predict the products of reactions, based on the classification of inorganic chemical reactions into four general categories: combination, decomposition, single-displacement, and double-displacement reactions.

Reactions may also be classified according to whether the oxidation number of one or more elements changes. Those reactions in which a change in oxidation number occurs are called oxidation–reduction reactions . One element increases its oxidation number (is oxidized), while the other decreases its oxidation number (is reduced).

Combination Reactions

In combination reactions, two substances, either elements or compounds, react to produce a single compound. One type of combination reaction involves two elements. Most metals react with most nonmetals to form ionic compounds. The products can be predicted from the charges expected for cations of the metal and anions of the nonmetal. For example, the product of the reaction between aluminum and bromine can be predicted from the following charges: 3+ for aluminum ion and 1− for bromide ion. Since there is a change in the oxidation numbers of the elements, this type of reaction is an oxidation–reduction reaction:

2Al (s ) + 3Br₂ (g ) → 2AlBr₃ (s )

Similarly, a nonmetal may react with a more reactive nonmetal to form a covalent compound. The composition of the product is predicted from the common oxidation numbers of the elements, positive for the less reactive and negative for the more reactive nonmetal (usually located closer to the upper right side of the Periodic Table). For example, sulfur reacts with oxygen gas to form gaseous sulfur dioxide:

S₈ (s ) + 8O₂ (g ) → 8SO₂ (g )

A compound and an element may unite to form another compound if in the original compound, the element with a positive oxidation number has an accessible higher oxidation number. Carbon monoxide, formed by the burning of hydrocarbons under conditions of oxygen deficiency, reacts with oxygen to form carbon dioxide:

2CO (g ) + O₂ (g ) → 2CO₂ (g )

The oxidation number of carbon changes from +2 to +4 so this reaction is an oxidation–reduction reaction.

Two compounds may react to form a new compound. For example, calcium oxide (or lime) reacts with carbon dioxide to form calcium carbonate (limestone):

CaO (s ) + CO₂ (g ) → CaCO₃ (s )

Decomposition Reactions

When a compound undergoes a decomposition reaction, usually when heated, it breaks down into its component elements or simpler compounds. The products of a decomposition reaction are determined largely by the identity of the anion in the compound. The ammonium ion also has characteristic decomposition reactions.

A few binary compounds decompose to their constituent elements upon heating. This is an oxidation–reduction reaction since the elements undergo a change in oxidation number. For example, the oxides and halides of noble metals (primarily Au, Pt, and Hg) decompose when heated. When red solid mercury(II) oxide is heated, it decomposes to liquid metallic mercury and oxygen gas:

2HgO (s ) → 2Hg (l ) + O₂ (g )

Some nonmetal oxides, such as the halogen oxides, also decompose upon heating:

2Cl₂O₅ (g ) → 2Cl₂ (g ) + 5O₂ (g )

Other nonmetal oxides, such as dinitrogen pentoxide, decompose to an element and a compound:

2N₂O₅ (g ) → O₂ (g ) + 4NO₂ (g )

Many metal salts containing oxoanions decompose upon heating. These salts either give off oxygen gas, forming a metal salt with a different nonmetal anion, or they give off a nonmetal oxide, forming a metal oxide. For example, metal nitrates containing Group 1A or 2A metals or aluminum decompose to metal nitrites and oxygen gas:

Mg(NO₃)₂ (s ) → Mg(NO₂)₂ (s ) + O₂ (g )

All other metal nitrates decompose to metal oxides, along with nitrogen dioxide and oxygen:

2Cu(NO₃)₂ (s ) → 2CuO (s ) + 4NO₂ (g ) + O₂ (g )

Salts of the halogen oxoanions decompose to halides and oxygen upon heating:

2KBrO₃ (s ) → 2KBr (s ) + 3O₂ (g )

Carbonates, except for those of the alkali metals, decompose to oxides and carbon dioxide.

CaCO₃ (s ) → CaO (s ) + CO₂ (g )

A number of compounds—hydrates, hydroxides, and oxoacids—that contain water or its components lose water when heated. Hydrates, compounds that contain water molecules, lose water to form anhydrous compounds, free of molecular water.

CaSO₄ · 2H₂O (s ) → CaSO₄ (s ) + 2H₂O (g )

Metal hydroxides are converted to metal oxides by heating:

2Fe(OH)₃ (s ) → Fe₂O₃ (s ) + 3H₂O (g )

Most oxoacids lose water until no hydrogen remains, leaving a nonmetal oxide:

H₂SO₄ (l ) → H₂O (g ) + SO₃ (g )

Oxoanion salts that contain hydrogen ions break down into the corresponding oxoanion salts and oxoacids:

Ca(HSO₄)₂ (s ) → CaSO₄ (s ) + H₂SO₄ (l )

Finally, some ammonium salts undergo an oxidation–reduction reaction when heated. Common salts of this type are ammonium dichromate, ammonium permanganate, ammonium nitrate, and ammonium nitrite. When these salts decompose, they give off nitrogen gas and water.

(NH₄)₂Cr₂O₇ (s ) → Cr₂O₃ (s ) + 4H₂O (g ) + N₂ (g )

2NH₄NO₃ (s ) → 2N₂ (g ) + 4H₂O (g ) + O₂ (g )

Ammonium salts, which do not contain an oxidizing agent, lose ammonia gas upon heating:

(NH₄)₂SO₄ (s ) → 2NH₃ (g ) + H₂SO₄ (l )

Single-Displacement Reactions

In a single-displacement reaction, a free element displaces another element from a compound to produce a different compound and a different free element. A more active element displaces a less active element from its compounds. These are all oxidation–reduction reactions. An example is the thermite reaction between aluminum and iron(III) oxide:

2Al (s ) + Fe₂O₃ (s ) → Al₂O₃ (s ) + 2Fe (l )

The element displaced from the compound is always the more metallic element—the one nearer the bottom left of the Periodic Table. The displaced element need not always be a metal, however. Consider a common type of single-displacement reaction, the displacement of hydrogen from water or from acids by metals.

The very active metals react with water. For example, calcium reacts with water to form calcium hydroxide and hydrogen gas. Calcium metal has an oxidation number of 0, whereas Ca²⁺ in Ca(OH)₂ has an oxidation number of +2, so calcium is oxidized. Hydrogen's oxidation number changes from +1 to 0, so it is reduced.

Ca (s ) + 2H₂O (l ) → Ca(OH)₂ (aq ) + H₂ (g )

Some metals, such as magnesium, do not react with cold water, but react slowly with steam:

Mg (s ) + 2H₂O (g ) → Mg(OH)₂ (aq ) + H₂ (g )

Still less active metals, such as iron, do not react with water at all, but react with acids.

Fe (s ) + 2HCl (aq ) → FeCl₂ (aq ) + H₂ (g )

Metals that are even less active, such as copper, generally do not react with acids.

To determine which metals react with water or with acids, we can use an activity series (see Figure 1), a list of metals in order of decreasing activity. Elements at the top of the series react with cold water. Elements above hydrogen in the series react with acids; elements below hydrogen do not react to release hydrogen gas.

The displacement of hydrogen from water or acids is just one type of single-displacement reaction. Other elements can also be displaced from their compounds. For example, copper metal reduces aqueous solutions of ionic silver compounds, such as silver nitrate, to deposit silver metal. The copper is oxidized.

Cu (s ) + 2AgNO₃ (aq ) → Cu(NO₃)₂ (aq ) + 2Ag (s )

The activity series can be used to predict which single-displacement reactions will take place. The elemental metal produced is always lower in the activity series than the displacing element. Thus, iron could be displaced from FeCl₂ by zinc metal but not by tin.

Double-Displacement Reactions

Aqueous barium chloride reacts with sulfuric acid to form solid barium sulfate and hydrochloric acid:

BaCl₂ (aq ) + H₂SO₄ (aq ) → BaSO₄ (s ) + 2HCl (aq )

Sodium sulfide reacts with hydrochloric acid to form sodium chloride and hydrogen sulfide gas:

Na₂S (aq ) + 2HCl (aq ) → 2NaCl (aq ) + H₂S (g )

Potassium hydroxide reacts with nitric acid to form water and potassium nitrate:

KOH (aq ) + HNO₃ (aq ) → H₂O (l ) + KNO₃ (aq )

These double-displacement reactions have two major features in common. First, two compounds exchange ions or elements to form new compounds. Second, one of the products is either a compound that will separate from the reaction mixture in some way (commonly as a solid or gas) or a stable covalent compound, often water.

Double-displacement reactions can be further classified as precipitation, gas formation, and acid–base neutralization reactions.

Precipitation Reactions

Precipitation reactions are those in which the reactants exchange ions to form an insoluble salt—one which does not dissolve in water. Reaction occurs when two ions combine to form an insoluble solid or precipitate. We predict whether such a compound can be formed by consulting solubility rules (see Table 1). If a possible product is insoluble, a precipitation reaction should occur.

A mixture of aqueous solutions of barium chloride and sodium sulfate contains the following ions: Ba²⁺ (aq ), Cl⁻ (aq ), Na⁺ (aq ), and SO₄²⁻ (aq ). According to solubility rules, most sulfate, sodium, and chloride salts are soluble. However, barium sulfate is insoluble. Since a barium ion and sulfate ion could combine to form insoluble barium sulfate, a reaction occurs.

BaCl₂ (aq ) + Na₂SO₄ (aq ) → BaSO₄ (s ) + 2NaCl (aq )

Gas-Formation Reactions

A double-displacement reaction should also occur if an insoluble gas is formed. All gases are soluble in water to some extent, but only a few gases [HCl (g ) and NH₃ (g )] are highly soluble. All other gases, generally binary covalent compounds, are sufficiently insoluble to provide a driving force if they are formed as a reaction product. For example, many sulfide salts will react with acids to form gaseous hydrogen sulfide:

ZnS (s ) + 2HCl (aq ) → ZnCl₂ (aq ) + H₂S (g )

Insoluble gases are often formed by the breakdown of an unstable double-displacement reaction product. For example, carbonates react with acids to form carbonic acid (H₂CO₃), an unstable substance that readily decomposes into water and carbon dioxide. Calcium carbonate reacts with hydrochloric acid to form calcium chloride and carbonic acid:

CaCO₃ (s ) + 2HCl (aq ) → CaCl₂ (aq ) + H₂CO₃ (aq )

Carbonic acid decomposes into water and carbon dioxide:

H₂CO₃ (aq ) → H₂O (l ) + CO₂ (g )

The net reaction is:

CaCO₃ (s ) + 2HCl (aq) → CaCl₂ (aq ) + H₂O (l ) + CO₂ (g )

Sulfites react with acids in a similar manner to release sulfur dioxide.

Acid-Base Neutralization Reactions

A neutralization reaction is a double-displacement reaction of an acid and a base. Acids are compounds that can release hydrogen ions; bases are compounds that can neutralize acids by reacting with hydrogen ions. The most common bases are hydroxide and oxide compounds of the metals. Normally, an acid reacts with a base to form a salt and water. Neutralization reactions occur because of the formation of the very stable covalent water molecule, H₂O, from hydrogen and hydroxide ions.

HCl (aq ) + NaOH (aq ) → NaCl (aq ) + H₂O (l )

Recognizing the pattern of reactants (element or compound, and the number of each) allows us to assign a possible reaction to one of the described classes. Recognizing the class of reaction allows us to predict possible products with some reliability.

see also Acid-Base Chemistry; Solution Chemistry; Thermodynamics.

James P. Birk

Bibliography

Adams, David L. (1999). "Issues-Directed Chemistry: Teaching Chemical Reactions Using Waste Treatment." Journal of Chemical Education 76:1088–1091.

Basolo, Fred (1980). "Systematic Inorganic Reaction Chemistry." Journal of Chemical Education 57:761–762.

Basolo, Fred (1984). "Teaching of Chemical Reactions and Synthesis." Journal of Chemical Education 61:520–521.

Basolo, Fred; and Parry, Robert W. (1980). "An Approach to Teaching Systematic Inorganic Reaction Chemistry in Beginning Chemistry Courses." Journal of Chemical Education 57:772–777.

Bent, Henry A., and Bent, Brian E. (1987). "Descriptive Chemistry." Journal of Chemical Education 64:249–251.

Cassen, T., and DuBois, Thomas D. (1982). "A Unified Approach to the Study of Chemical Reactions in Freshman Chemistry." Journal of Chemical Education 59: 377–379.

Charola, A. Elena (1987). "Acid Rain: Effect on Stone Monuments." Journal of Chemical Education 64:436–437.

Lambda, Ram S., Sharma, Shiva, and Lloyd, Baird W. (1997). "Constructing Chemical Concepts through a Study of Metals and Metal Ions." Journal of Chemical Education 74:1095–1099.

Ragsdale, Ronald O., and Zipp, Arden P. (1992). "Helping Students to Improve Their Approach to Predicting the Products of Chemical Reactions." Journal of Chemical Education 69:390–392.

Robson, David P. (1987). "Sunken Treasure." ChemMatters 5(2):4–9.

Internet Resources

"Balancing Equations & Reaction Types." Chem Team. Available from <http://dbhs.wvusd.k12.ca.us/Equations/Equations.html>.

ChemWeb 2000. Available from <http://library.thinkquest.org/19957/>.

Chemical Reactions

KEY TERMS

Resources

Chemical reactions describe the changes between reactants (the initial substances that enter into the reaction) and products (the final substances that are present at the end of the reaction). Chemical reactions involve a rearrangement of the atoms in reactants to form products with new structures in such a way as to conserve atoms. Chemical equations are notations that are used to summarize and convey information regarding chemical reactions.

In a balanced chemical reaction, all of the matter (i.e., atoms or molecules) that enters into a reaction must be accounted for in the products of a reaction. Accordingly, associated with the symbols for the reactants and products are numbers (stoichiometry coefficients) that represent the number of molecules, formula units, or moles of a particular reactant or product. Reactants and products are separated by addition symbols (plus signs). These represent the interaction of the reactants and are used to separate and list the products formed. The chemical equations for some reactions may have a lone reactant or a single product. The subscript numbers associated with the chemical formula designating individual reactants and products represent the number of atoms of each element that are in each molecule (for covalently-bonded substances) or formula unit (for ionically-associated substances) of reactants or products.

For a chemical reaction to be balanced, all of the atoms present in molecules, formula units, or moles of reactants to the left of the equation arrow must be present in the molecules, formula units, or moles of product to the right of the equation arrow. The combinations of the atoms may change (indeed, this is what chemical reactions do) but the number of atoms present in reactants must equal the number of atoms present in products. Electrical charge is also conserved between reactants and products in balanced chemical reactions.

Although chemical equations are usually concerned only with reactants and products, chemical reactions may proceed through multiple intermediate steps. In such multistep reactions, the products of one become the reactants (intermediary products) for the next step in the sequence.

Reaction catalysts are chemical species that alter the energy requirements of reactions and thereby alter the speed at which reactions run (i.e., they control the rate of formation of products).

Combustion reactions are those in which oxygen combines with another compound to form water and carbon dioxide. The equations for these reactions usually designate that the reaction is exothermic (heat producing). Synthesis reactions occur when two or more simple compounds combine to form a more complicated compound. Decomposition reactions reflect the reversal of synthesis reactions (e.g., reactions where complex molecules are broken down into simpler molecules). The electrolysis of water to make oxygen and hydrogen is an excellent example of a decomposition reaction.

Equations for single-displacement reactions, double-displacement, and acid-base reactions reflect the appropriate reallocation of atoms in the products.

In accord with the laws of thermodynamics, all chemical reactions change the energy state of the reactants. The change in energy results from changes in the number and strengths of chemical bonds as the reaction proceeds. The heat of reaction is defined as the quantity of heat evolved or absorbed during a chemical reaction. A reaction is called exothermic if heat is released or given off during a chemical transformation. Alternatively, in an endothermic reaction, heat is absorbed in transforming reactants into products. In endothermic reactions, heat energy must be supplied to the system for a reaction to occur and the heat content of the products is larger than that of the reactants. For example, if a mixture of gaseous hydrogen and oxygen is ignited, water is formed and heat energy is given off. The chemical reaction is an exothermic reaction and the heat content of the product(s) is lower than that for the reactants. The study of energy utilization in chemical reactions is called chemical kinetics and is important in understanding chemical transformations.

A chemical reaction takes place in a vessel that can be treated as a system. If the heat “flows” into the vessel during reaction, the reaction is said to be “endothermic” (e.g., a decomposition process) and the amount of heat, say, q, provided to the system is taken as a positive quantity. On the other hand, when the system has lost heat to the outside world, the reaction is “exothermic” (e.g., a combustion process) and q is viewed as a negative number. Normally the heat change involved in a reaction can be measured in an adiabatic bomb calorimeter. The reaction is initiated inside a constant-volume container. The observed change in temperature and the information on the total heat capacity of the colorimeter are employed to calculate q. If the heat of reaction is obtained for both the products and reactants at the same temperature after reaction and also in their standard states, it is then defined as the “standard heat of reaction,” denoted by ΔH°.

Both chemical kinetics and thermodynamics are crucial issues in studying chemical reactions. Chemical kinetics help us search for the factors that influence the rate of reaction. It tells us how fast the chemical reaction will take place and about what the sequence of individual chemical events is to produce observed reactions. Very often, a single reaction like A → B may take several steps to complete. In other words, a chain reaction mechanism, which can include initiation, propagation, and termination stages, is involved, and the individual reaction rates may be very different. With a search for actual reaction mechanisms, the expression for overall reaction rate can be given correctly. As to determining the maximum extent to which a chemical reaction can proceed, and how much heat will be absorbed or liberated, we need to estimate from thermodynamics data. Therefore, kinetic and thermodynamic information is extremely important for reactor design.

As an example of a chemical reaction, hydrogen (H₂ ) and oxygen (O₂ ) gases under certain conditions can react to form water (H₂ O). Water then exists as solid (ice), liquid, or vapor (steam); they all have the same composition, H₂ O, but exhibit a difference in how H₂ O molecules are brought together due to variations in temperature and pressure.

Chemical reactions can take place in one phase alone and are termed “homogeneous.” They can also proceed in the presence of at least two phases, such as reduction of iron ore to iron and steel, which are normally described as “heterogeneous” reactions. Quite frequently, the rate of chemical reaction is altered by foreign materials, so-called catalysts, that are neither reactants nor products. Although usually used to accelerate reactions, reaction catalysts can either accelerate or hinder the reaction process. Typical examples are found in Pt as the catalyst for oxidation of sulfur dioxide (SO₂ ) and iron promoted with Al₂ O₃ and K as the catalyst for ammonia (NH₃ ) synthesis.

Chemical reactions are characterized as irreversible, reversible, or oscillating. In the former case, the equilibrium for the reaction highly favors formation of the products, and only a very small amount of reactants remains in the system at equilibrium. In contrast to this, a reversible reaction allows for appreciable quantities of all reactants and products co-existing at equilibrium. H₂ O←+ 3NO₂ ⇄ 2HNO₃ ←+ NO is an

KEY TERMS

Chemical kinetics —The study of the reaction mechanism and rate by which one chemical species is converted to another.

Equilibrium —The conditions under which a system shows no tendency for a change in its state. At equilibrium, the net rate of reaction becomes zero.

Phase —A homogeneous region of matter.

Standard state —The state defined in reaction thermodynamics for calculation purposes in which the pure gas in the ideal-gas state at 1 atm and pure liquid or solid at 1 atm are taken for gas and liquid or solid, respectively.

Thermodynamics —Thermodynamics is the study of energy in the form of heat and work, and the relationship between the two.

example of a reversible reaction. In an oscillating chemical reaction, the concentrations of the reactants and products change with time in a periodic or quasi-periodic manner. Chemical oscillators exhibit chaotic behavior in which concentrations of products and the course of a reaction depend on the initial conditions of the reaction.

Chemical reactions may proceed in various ways: a single reaction A → B; series reactions A→ B→ C; side-by-side parallel reactions A→ B and C→ D; two competitive parallel reactions A→ B and A→ C; or mixed parallel and series reactions A+ B→ C and C+ B→ D. In order for chemical reactions to occur, reactive species have to first encounter each other so that they can exchange atoms or groups of atoms. In gas phases, this step relies on collision, whereas in liquid and solid phases, diffusion process (mass transfer) plays a key role. However, even reactive species do encounter each other, and certain energy state changes are required to surmount the energy barrier for the reaction. Normally, this minimum energy requirement (e.g., used to break old chemical bonds and to form new ones) varies with temperature, pressure, the use of catalysts, etc. In other words, the rate of chemical reaction depends heavily on encounter rates or frequencies and energy availability, and it can vary from a value approaching infinity to essentially zero.

See also Catalyst and catalysis; Chemical bond; Chemistry; Conservation laws; Entropy; Enzyme; Equation, chemical; Equilibrium, chemical; Molecular formula; Moles; Stereochemistry.

Resources

BOOKS

Housecroft, Catherine E., and Alan G. Sharpe. Inorganic Chemistry, 2nd ed. Upper Saddle River, N.J.: Prentice Hall, 2005.

Incropera, Frank P., and David P. DeWitt. Fundamentals of Heat and Mass Transfer, 5th ed. New York: John Wiley & Sons, 2001.

Moran, Michael J., and Howard N. Shapiro. Fundamentals of Engineering Thermodynamics, 4th ed. New York: John Wiley & Sons, 2000.

OTHER

Carpi, Anthony. “Chemical Reactions” City University of New York, John Jay College of Criminal Justice <http://web.jjay.cuny.edu/äcarpi/NSC/6-react.htm> (accessed November 13, 2006).

Carpi, Anthony. “Chemical Reactions” Visionlearning

<http://www.visionlearning.com/library/module_viewer.php?mid=54> (accessed November 13, 2006).

K. Lee Lerner

Pang-Jen Kung

Chemical reactions

Chemical reactions describe the changes between reactants (the initial substances that enter into the reaction) and products (the final substances that are present at the end of the reaction). Describing interactions among chemical species, chemical reactions involve a rearrangement of the atoms in reactants to form products with new structures in such a way as to conserve atoms. Chemical equations are notations that are used to concisely summarize and convey information regarding chemical reactions.

In a balanced chemical reaction all of the matter (i.e., atoms or molecules) that enter into a reaction must be accounted for in the products of a reaction. Accordingly, associated with the symbols for the reactants and products are numbers (stoichiometry coefficients) that represent the number of molecules, formula units, or moles of a particular reactant or product. Reactants and products are separated by addition symbols (addition signs). The addition signs represent the interaction of the reactants and are used to separate and list the products formed. The chemical equations for some reactions may have a lone reactant or a single product. The subscript numbers associated with the chemical formula designating individual reactants and products represent the number of atoms of each element that are in each molecule (for covalently bonded substances) or formula unit (for ironically associated substances) of reactants or products.

For a chemical reaction to be balanced, all of the atoms present in molecules or formula units or moles of reactants to the left of the equation arrow must be present in the molecules, formula units and moles of product to the right of the equation arrow. The combinations of the atoms may change (indeed, this is what chemical reactions do) but the number of atoms present in reactants must equal the number of atoms present in products.

Charge is also conserved in balanced chemical reactions and therefore there is a conservation of electrical charge between reactants and products.

Although chemical equations are usually concerned only with reactants and products chemical reactions may proceed through multiple intermediate steps. In such multi-step reactions the products of one reaction become the reactants (intermediary products) for the next step in the reaction sequence.

Reaction catalysts are chemical species that alter the energy requirements of reactions and thereby alter the speed at which reactions run (i.e., control the rate of formation of products).

Combustion reactions are those where oxygen combines with another compound to form water and carbon dioxide . The equations for these reactions usually designate that the reaction is exothermic (heat producing). Synthesis reactions occur when two or more simple compounds combine to form a more complicated compound. Decomposition reactions reflect the reversal of synthesis reactions (e.g., reactions where complex molecules are broken down into simpler molecules). The electrolysis of water to make oxygen and hydrogen is an excellent example of a decomposition reaction.

Equations for single displacement reactions, double displacement, and acid-base reactions reflect the appropriate reallocation of atoms in the products.

In accord with the laws of thermodynamics , all chemical reactions change the energy state of the reactants. The change in energy results from changes in the in the number and strengths of chemical bonds as the reaction proceeds. The heat of reaction is defined as the quantity of heat evolved or absorbed during a chemical reaction. A reaction is called exothermic if heat is released or given off during a chemical transformation. Alternatively, in an endothermic reaction, heat is absorbed in transforming reactants to products. In endothermic reactions, heat energy must be supplied to the system in order for a reaction to occur and the heat content of the products is larger than that of the reactants. For example, if a mixture of gaseous hydrogen and oxygen is ignited, water is formed and heat energy is given off. The chemical reaction is an exothermic reaction and the heat content of the product(s) is lower than that for the reactants. The study of energy utilization in chemical reactions is called chemical kinetics and is important in understanding chemical transformations.

A chemical reaction takes place in a vessel which can be treated as a system. If the heat "flows" into the vessel during reaction, the reaction is said to be "endothermic" (e.g., a decomposition process) and the amount of heat, say, q, provided to the system is taken as a positive quantity. On the other hand, when the system has lost heat to the outside world, the reaction is "exothermic" (e.g., a combustion process) and q is viewed as a negative number. Normally the heat change involved in a reaction can be measured in an adiabatic bomb calorimeter. The reaction is initiated inside a constant-volume container. The observed change in temperature and the information on the total heat capacity of the colorimeter are employed to calculate q. If the heat of reaction is obtained for both the products and reactants at the same temperature after reaction and also in their standard states, it is then defined as the "standard heat of reaction," denoted by ΔH°.

Both chemical kinetics and thermodynamics are crucial issues in studying chemical reactions. Chemical kinetics help us search for the factors that influence the rate of reaction. It provides us with the information about how fast the chemical reaction will take place and about what the sequence of individual chemical events is to produce observed reactions. Very often, a single reaction like A B may take several steps to complete. In other words, a chain reaction mechanism is actually involved which can include initiation, propagation, and termination stages, and their individual reaction rates may be very different. With a search for actual reaction mechanisms, the expression for overall reaction rate can be given correctly. As to determining the maximum extent to which a chemical reaction can proceed and how much heat will be absorbed or liberated, we need to estimate from thermodynamics data. Therefore, kinetic and thermodynamic information is extremely important for reactor design.

As an example of a chemical reaction, Hydrogen (H2) and oxygen (O2) gases under certain conditions can react to form water (H2O). Water then exists as solid (ice ), liquid, or vapor (steam); they all have the same composition, H2O, but exhibit a difference in how H2O molecules are brought together due to variations in temperature and pressure .

Chemical reactions can take place in one phase alone and are termed "homogeneous." They can also proceed in the presence of at least two phases, such as reduction of iron ore to iron and steel , which are normally described as "heterogeneous" reactions. Quite frequently, the rate of chemical reaction is altered by foreign materials, so-called catalysts, that are neither reactants nor products. Although usually used to accelerate reactions, reaction catalysts can either accelerate or hinder the reaction process. Typical examples are found in Pt as the catalyst for oxidation of sulfur dioxide (SO2) and iron promoted with Al2O3 and K as the catalyst for ammonia (NH3) synthesis.

Chemical reactions can be characterized as irreversible, reversible, or oscillating. In the former case, the equilibrium for the reaction highly favors formation of the products, and only a very small amount of reactants remains in the system at equilibrium. In contrast to this, a reversible reaction allows for appreciable quantities of all reactants and products co-existing at equilibrium. H2O + 3NO2 ⇄ 2HNO3 + NO is an example of a reversible reaction. In an oscillating chemical reaction, the concentrations of the reactants and products change with time in a periodic or quasi-periodic manner. Chemical oscillators exhibit chaotic behavior, in which concentrations of products and the course of a reaction depend on the initial conditions of the reaction.

Chemical reactions may proceed as a single reaction A → B, series reactions A → B → C, side-by-side parallel reactions A → B and C → D, two competitive parallel reactions B and A C, or mixed parallel and series reactions A + B → C and C + B → D. In order for chemical reactions to occur, reactive species have to first encounter each other so that they can exchange atoms or groups of atoms. In gas phases, this step relies on collision, whereas in liquid and solid phases, diffusion process (mass transfer) plays a key role. However, even reactive species do encounter each other, and certain energy state changes are required to surmount the energy barrier for the reaction. Normally, this minimum energy requirement (e.g., used to break old chemical bonds and to form new ones) is varied with temperature, pressure, the use of catalysts, etc. In other words, the rate of chemical reaction depends heavily on encounter rates or frequencies and energy availability, and it can vary from a value approaching infinity to essentially zero .

See also Catalyst and catalysis; Chemical bond; Chemistry; Conservation laws; Entropy; Enzyme; Equation, chemical; Equilibrium, chemical; Molecular formula; Moles; Stereochemistry.

Resources

books

Housecroft, Catherine E., et al. Inorganic Chemistry. Prentice Hall, 2001.

Incropera, Frank P., and David P. DeWitt. Fundamentals of Heat and Mass Transfer. 5th ed. John Wiley & Sons, 2001.

Moran, Michael J., and Howard N. Shapiro. Fundamentals of Engineering Thermodynamics. 4th ed. John Wiley & Sons, 2000.

K. Lee Lerner Pang-Jen Kung

KEY TERMS

Chemical kinetics

—The study of the reaction mechanism and rate by which one chemical species is converted to another.

Equilibrium

—The conditions under which a system shows no tendency for a change in its state. At equilibrium the net rate of reaction becomes zero.

Phase

—A homogeneous region of matter.

Standard state

—The state defined in reaction thermodynamics for calculation purposes in which the pure gas in the ideal-gas state at 1 atm and pure liquid or solid at 1 atm are taken for gas and liquid or solid, respectively.

Thermodynamics

—Thermodynamics is the study of energy in the form of heat and work, and the relationship between the two.

Reaction, chemical

When a chemical reaction occurs, at least one product is formed that is different from the substances present before the change occurred. As an example, it is possible to pass an electric current through a sample of water and obtain a mixture of oxygen and hydrogen gases. That change is a chemical reaction because neither oxygen nor hydrogen were present as elements before the change took place.

Any chemical change involves two sets of substances: reactants and products. A reactant is an element or compound present before a chemical change takes place. In the example above, only one reactant was present: water. A product is an element or compound formed as a result of the chemical reaction. In the preceding example, both hydrogen and oxygen are products of the reaction.

Chemical reactions are represented by means of chemical equations. A chemical equation is a symbolic statement that represents the changes that occur during a chemical reaction. The statement consists of the symbols of the elements and the formulas of the products and reactants, along with other symbols that represent certain conditions present in the reaction. For example, the arrow (or yields) sign, *, separates the reactants from the products in a reaction. The chemical equation that represents the electrolysis of water is 2 H₂O → 2 H₂ + O₂.

Types of chemical reactions

Most chemical reactions can be categorized into one of about five general types: synthesis, decomposition, single replacement, double replacement, and oxidation-reduction. A miscellaneous category is also needed for reactions that do not fit into one of these five categories.

Characteristics of each type.

Synthesis: Two substances combine to form one new substance:

In general: A + B → AB

For example:

2 Na + Cl₂ → 2 NaCl or CaO + H₂O → Ca(OH)₂

Decomposition: One substance breaks down to form two new substances:

In general: AB → A + B

For example: 2 H₂O → 2 H₂ + O₂

Single Replacement: An element and a compound react such that the element replaces one other element in the compound:

In general: A + BC → AC + B

For example: Mg + 2 HCl → MgCl₂ + H₂

Double Replacement: Two compounds react with each other in such a way that they exchange partners with each other:

In general: AB + CD → AD + CB

For example:

NaBr + HCl → NaCl + HBr

Oxidation-reduction: One or more elements in the reaction changes its oxidation state during the reaction: In general: A³⁺ → A⁶⁺

For example: Cr³⁺ → Cr⁶⁺

Energy changes and chemical kinetics

Chemical reactions are typically accompanied by energy changes. The equation for the synthesis of ammonia from its elements is N₂ + 3 H₂ → 2 NH₃, but that reaction takes place only under very special conditions—namely at a high temperature and pressure and in the presence of a catalyst. Energy changes that occur during chemical reactions are the subject of a field of science known as thermodynamics.

In addition, chemical reactions are often a good deal more complex than a chemical equation might lead one to believe. For example, one can write the equation for the synthesis of hydrogen iodide from its elements, as follows: H₂ + I₂ → 2 HI. In fact, chemists know that this reaction does not take place in a single step. Instead, it occurs in a series of reactions in which hydrogen and iodine atoms react with each other one at a time. The final equation, H₂ + I₂ → 2 HI, is actually no more than a summary of the net result of all those reactions. The field of chemistry that deals with the details of chemical reactions is known as chemical kinetics.