Reactions are the "verbs" of chemistry—the activity that chemists study. Many reactions move to their conclusion and then stop, meaning that the reactants have been completely transformed into products, with no means of returning to their original state. In some cases, the reaction truly is irreversible, as for instance when combustion changes both the physical and chemical properties of a substance. There are plenty of other circumstances, however, in which a reverse reaction is not only possible but an ongoing process, as the products of the first reaction become the reactants in a second one. This dynamic state, in which the concentrations of reactants and products remains constant, is referred to as equilibrium. It is possible to predict the behavior of substances in equilibrium through the use of certain laws, which are applied in industries seeking to lower the costs of producing specific chemicals. Equilibrium is also useful in understanding processes that preserve—or potentially threaten—human health.
HOW IT WORKS
Chemical Reactions in Brief
What follows is a highly condensed discussion of chemical reactions, and particularly the methods for writing equations to describe them. For a more detailed explanation of these principles, the reader is encouraged to consult the Chemical Reactions essay.
A chemical reaction is a process whereby the chemical properties of a substance are altered by a rearrangement of the atoms in the substance. The changes produced by a chemical reaction are fundamentally different from physical changes, such as boiling or melting liquid water, changes that alter the physical properties of water without affecting its molecular structure.
INDICATIONS THAT A CHEMICAL REACTION HAS OCCURRED.
Though chemical reactions are most effectively analyzed in terms of molecular properties and behaviors, there are numerous indicators that suggest to us when a chemical reaction has occurred. It is unlikely that all of these will result from any one reaction, and in fact chances are that a particular reaction will manifest only one or two of these effects. Nonetheless, these offer us hints that a reaction has taken place:
Signs that a substance has undergone a chemical reaction:
- Water is produced
- A solid forms
- Gases are produced
- Bubbles are formed
- There is a change in color
- The temperature changes
- The taste of a consumable substance changes
- The smell changes
CHEMICAL CHANGES CONTRASTED WITH PHYSICAL CHANGES OF TEMPERATURE.
Many of these effects can be produced simply by changing the temperature of a substance, but again, the mere act of applying heat from outside (or removing heat from the substance itself) does not constitute a chemical change. Water can be "produced" by melting ice, but the water was already there—it only changed form. By contrast, when an acid and a base react to form water and a salt, that is a true chemical reaction.
Similarly, the freezing of water forms a solid, but no new substance has been formed; in a chemical reaction, by contrast, two liquids can react to form a solid. When water boils through the application of heat, bubbles form, and a gas or vapor is produced; yet in chemical changes, these effects are not the direct result of applying heat.
In this context, a change in temperature, noted as another sign that a reaction has taken place, is a change of temperature from within the substance itself. Chemical reactions can be classified as heat-producing (exothermic) or heat-absorbing (endothermic). In either case, the transfer of heat is not accomplished simply by creating a temperature differential, as would occur if heat were transferred merely through physical means.
Why Do Chemical Reactions Occur?
At one time, chemists could only study reactions from "outside," as it were—purely in terms of effects noticeable through the senses. Between the early nineteenth and the early twentieth centuries, however, the entire character of chemistry changed, as did the terms in which chemists discussed reactions. Today, those reactions are analyzed primarily in terms of subatomic, atomic, and molecular properties and activities.
Despite all this progress, however, chemists still do not know exactly what happens in a chemical reaction—but they do have a good approximation. This is the collision model, which explains chemical reactions in terms of collisions between molecules. If the collision is strong enough, it can break the chemical bonds in the reactants, resulting in a re-formation of atoms within different molecules. The more the molecules collide, the more bonds are being broken, and 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. By raising the temperature, the speeds of the molecules themselves increase, and the collisions possess more energy. A certain energy threshold, the activation energy (symbolized E a) must be crossed in order for a reaction to occur. A temperature increase raises the likelihood that a given collision will cross the activation-energy threshold, producing the energy to break the molecular bonds and promote the chemical reaction.
Raising the temperature and the concentrations of reactants can increase the energy and hasten the reactions, but in some cases it is not possible to do either. Fortunately, the rate of reaction can be increased in a third way, through the introduction 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. Many chemistry textbooks discuss catalysts within the context of equilibrium; however, because catalysts play such an important role in human life, in this book they are the subject of a separate essay.
Chemical Equations Involving Equilibrium
A chemical equation, like a mathematical equation, symbolizes an interaction between entities that produces a particular result. In the case of a chemical equation, the entities are not numbers but reactants, and they interact with each other not through addition or multiplication, but by chemical reaction. Yet just as a product is the result of multiplication in mathematics, a product in a chemical equation is the substance or substances that result from the reaction.
Instead of an equals sign between the reactants and the product, an arrow is used. When the arrow points to the right, this indicates a forward reaction; conversely, an arrow pointing to the left symbolizes a reverse reaction. In a reverse reaction, the products of a forward reaction have become the reactants, and the reactants of the forward reaction are now the products. This is indicated by an arrow that points toward the left.
Chemical equilibrium, which occurs when the ratio between the reactants and products is constant, and in which the forward and reverse reactions take place at the same rate, is symbolized thus: ⇌. Note that the arrows, the upper one pointing right and the lower one pointing left, are of the same length. There may be certain cases, discussed below, in which it is necessary to show these arrows as unequal in length as a means of indicating the dominance of either the forward or reverse reaction.
Chemical equations usually include notation indicating the state or phase of matter for the reactants and products: (s) for a solid; (l) for a liquid; (g) for a gas. A fourth symbol, (aq), indicates a substance dissolved in water—that is, an aqueous solution. In the following paragraphs, we will apply a chemical equation to the demonstration of equilibrium, but will not discuss the balancing of equations. The reader is encouraged to consult the passage in the Chemical Reactions essay that addresses that process, vital to the recording of accurate data.
A SIMPLE EQUILIBRIUM EQUATION.
Let us now consider a simple equation, involving the reaction between water and carbon monoxide (CO) at high temperatures in a closed container. The initial equation is written thus: H2O(g) + CO(g) →H2(g) + CO2g). In plain English, water in the gas phase (steam) has reacted with carbon monoxide to produce hydrogen gas and carbon dioxide.
As the reaction proceeds, the amount of reactants decreases, and the concentration of products increases. At some point, however, there will be a balance between the numbers of products and reactants—a state of chemical equilibrium represented by changing the right-pointing arrow to an equilibrium symbol: H2O(g) + CO(g) ⇋ H2(g) + CO2g). Assuming that the system is not disturbed (that is, that the container is kept closed and no outside substances are introduced), equilibrium will continue to be maintained, because the reverse reaction is occurring at the same rate as the forward one.
Note what has been said here: reactions are still occurring, but the forward and rearward reactions balance one another. Thus equilibrium is not a static condition, but a dynamic one, and indeed, chemical equilibrium is sometimes referred to as "dynamic equilibrium." On the other hand, some chemists refer to chemical equilibrium simply as equilibrium, but here the qualifier chemical has been used to distinguish this from the type of equilibrium studied in physics. Physical equilibrium, which involves factors such as center of gravity, does help us to understand chemical equilibrium, but it is a different phenomenon.
Homogeneous and Heterogeneous Equilibria
It should be noted that the equation used above identifies a situation of homogeneous equilibrium, in which all the substances are in the same phase or state of matter—gas, in this case. It is also possible to achieve chemical equilibrium in a reaction involving substances in more than one phase of matter.
An example of such heterogeneous equilibrium is the decomposition of calcium carbonate for the production of lime, a process that involves the application of heat. Here the equation would be written thus: CaCO3(s) ⇋ CaO(s) + CO2(g). Both the calcium carbonate (CaCO3) and the lime (CaO) are solids, whereas the carbon dioxide produced in this reaction is a gas.
The Equilibrium Constant
In 1863, Norwegian chemists Cato Maximilian Guldberg (1836-1902) and Peter Waage (1833-1900)—who happened to be brothers-in-law—formulated what they called the law of mass action. Today, this is called the law of chemical equilibrium, which states that the direction taken by a reaction is dependant not merely on the mass of the various components of the reaction, but also upon the concentration—that is, the mass present in a given volume.
This can be expressed by the formula a A + b B ⇌ c C + d D, where the capital letters represent chemical species, and the italicized lowercase letters indicate their coefficients. The equation [C] c[D]d/[A] a[B]b yields what is called an equilibrium constant, symbolized K.
The above formula expresses the equilibrium constant in terms of molarity, the amount of solute in a given volume of solution, but in the case of gaseous reactants and products, the equilibrium constant can also be expressed in terms of partial pressures. In the reaction of water and carbon monoxide to produce hydrogen molecules and carbon dioxide (H2O + CO ⇋ H2 + CO2). In chemical reactions involving solids, however, the concentration of the solid—because it is considered to be invariant—does not appear in the equilibrium constant. In the reaction described earlier, in which calcium carbonate was in equilibrium with solid lime and gaseous carbon dioxide, K = pressure of CO2.
We will not attempt here to explore the equilibrium constant in any depth, but it is important to recognize its usefulness. For a particular reaction at a specific temperature, the ratio of concentrations between reactants and products will always have the same value—the equilibrium constant, or K. Because it is not dependant on the amounts of reactants and products mixed together initially, K remains the same: the concentrations themselves may vary, but the ratios between the concentrations in a given situation do not.
Le ChÂtelier's Principle
Not all situations of equilibrium are alike: depending on certain factors, the position of equilibrium may favor one side of the equation or the other. If a company is producing chemicals for sale, for example, its production managers will attempt to influence reactions in such a way as to favor the forward reaction. In such a situation, it is said that the equilibrium position has been shifted to the right. In terms of physical equilibrium, mentioned above, this would be analogous to what would happen if you were holding your arms out on either side of your body, with a heavy lead weight in your left hand and a much smaller weight in the right hand.
Your center of gravity, or equilibrium position, would shift to the left to account for the greater force exerted by the heavier weight.
A value of K significantly above 1 causes a shift to the right, meaning that at equilibrium, there will be more products than reactants. This is a situation favorable to a chemical company's managers, who desire to create more of the product from less of the reactants. However, nature abhors an imbalance, as expressed in Le Châtelier's principle. Named after French chemist Henri Le Châtelier (1850-1936), this principle maintains that whenever a stress or change is imposed on a chemical system in equilibrium, the system will adjust the amounts of the various substances to reduce the impact of that stress.
Suppose we add more of a particular substance to increase the rate of the forward reaction. In an equation for this reaction, the equilibrium symbol is altered, with a longer arrow pointing to the right to indicate that the forward reaction is favored. Again, the equilibrium position has shifted to the right—just as one makes physical adjustments to account for an imbalanced weight. The system responds by working to consume more of the reactant, thus adjusting to the stress that was placed on it by the addition of more of that substance. By the same token, if we were to remove a particular reactant or product, the system would shift in the direction of the detached component.
Note that Le Châtelier's principle is mathematically related to the equilibrium constant. Suppose we have a basic equilibrium equation of A + B ⇌ C, with A and B each having molarities of 1, and C a molarity of 4. This tells us that K is equal to the molarity of C divided by that of A multiplied by B = 4/(1 · 1). Suppose, now, that enough of C were added to bring its concentration up to 6. This would mean that the system was no longer at equilibrium, because C/(A · B) no longer equals 4. In order to return the ratio to 4, the numerator (C) must be decreased, while the denominator (A · B) is increased. The reaction thus shifts from right to left.
CHANGES IN VOLUME AND TEMPERATURE.
If the volume of gases in a closed container is decreased, the pressure increases. An equilibrium system will therefore shift in the direction that reduces the pressure; but if the volume is increased, thus reducing the pressure, the system will respond by shifting to increase pressure. Note, however, that not all increases in pressure lead to a shift in the equilibrium. If the pressure were increased by the addition of a noble gas, the gas itself—since these elements are noted for their lack of reactivity—would not be part of the reaction. Thus the species added would not be part of the equilibrium constant expression, and there would be no change in the equilibrium.
In any case, no change in volume alters the equilibrium constant K ; but where changes in temperature are involved, K is indeed altered. In an exothermic, or heat-producing reaction, the heat is treated as a product. Thus, when nitrogen and hydrogen react, they produce not only ammonia, but a certain quantity of heat. If this system is at equilibrium, Le Châtelier's principle shows that the addition of heat will induce a shift in equilibrium to the left—in the direction that consumes heat or energy.
The reverse is true in an endothermic, or heat-absorbing reaction. As in the process described earlier, the thermal decomposition of calcium carbonate produces lime and carbon dioxide. Because heat is used to cause this reaction, the amount of heat applied is treated as a reactant, and an increase in temperature will cause the equilibrium position to shift to the right.
Equilibrium and Health
Discussions of chemical equilibrium tend to be rather abstract, as the foregoing sections on the equilibrium constant and Le Châtelier's principle illustrate. (The reader is encouraged to consult additional sources on these topics, which involve a number of particulars that have been touched upon only briefly here.) Despite the challenges involved in addressing the subject of equilibrium, the results of chemical equilibrium can be seen in processes involving human health.
The cooling of food with refrigerators, along with means of food preservation that do not involve changes in temperature, maintains chemical equilibrium in the foods and thereby prevents or at least retards spoilage. Even more important is the maintenance of equilibrium in reactions between hemoglobin and oxygen in human blood.
HEMOGLOBIN AND OXYGEN.
Hemoglobin, a protein containing iron, is the material in red blood cells responsible for transporting oxygen to the cells. Each hemoglobin molecule attaches to four oxygen atoms, and the equilibrium conditions of the hemoglobin-oxygen interaction can be expressed thus: Hb(aq) + 4O2(g) ⇋ Hb(O2)4(aq), where "Hb" stands for hemoglobin. As long as there is sufficient oxygen in the air, a healthy equilibrium is maintained; but at high altitudes, considerable changes occur.
At significant elevations above sea level, the air pressure is lowered, and thus it is more difficult to obtain the oxygen one needs. The result, in accordance with Le Châtelier's principle, is a shift in equilibrium to the left, away from the oxygenated hemoglobin. Without adequate oxygen fed to the body's cells and tissues, a person tends to feel light-headed.
When someone not physically prepared for the change is exposed to high altitudes, it may be necessary to introduce pressurized oxygen from an oxygen tank. This shifts the equilibrium to the right. For people born and raised at high altitudes, however, the body's chemistry performs the equilibrium shift—by producing more hemoglobin, which also shifts equilibrium to the right.
HEMOGLOBIN AND CARBON MONOXIDE.
When someone is exposed to carbon monoxide gas, a frightening variation on the normal hemoglobin-oxygen interaction occurs. Carbon monoxide "fools" hemoglobin into mistaking it for oxygen because it also bonds to hemoglobin in groups of four, and the equilibrium expression thus becomes: Hb(aq) + 4CO(g) ⇋ Hb(CO)4(aq). Instead of hemoglobin, what has been produced is called carboxyhemoglobin, which is even redder than hemoglobin. Therefore, one sign of carbon monoxide poisoning is a flushed face.
The bonds between carbon monoxide and hemoglobin are about 300 times as strong as those between hemoglobin and oxygen, and this means a shift in equilibrium toward the right side of the equation—the carboxyhemoglobin side. It also means that K for the hemoglobin-carbon monoxide reaction is much higher than for the hemoglobin-oxygen reaction. Due to the affinity of hemoglobin for carbon monoxide, the hemoglobin puts a priority on carbon monoxide bonds, and hemoglobin that has bonded with carbon monoxide is no longer available to carry oxygen.
Carbon monoxide in small quantities can cause headaches and dizziness, but larger concentrations can be fatal. To reverse the effects of the carbon monoxide, pure oxygen must be introduced to the body. It will react with the carboxyhemoglobin to produce properly oxygenated hemoglobin, along with carbon monoxide: Hb(CO)4(aq) + 4O2(g) ⇋ Hb(O2)4(aq) + 4CO(g). The gaseous carbon monoxide thus produced is dissipated when the person exhales.
WHERE TO LEARN MORE
"Catalysts" (Web site). <http://edie.cprost.sfu.ca/~rhlogan/catalyst.html> (June 9, 2001).
Challoner, Jack. The Visual Dictionary of Chemistry. New York: DK Publishing, 1996.
"Chemical Equilibrium." Davidson College Department of Chemistry (Web site). <http://www.chm.davidson.edu/ronutt/che115/EquKin.htm> (June 9, 2001).
"Chemical Equilibrium in the Gas Phase." Virginia Tech Chemistry Department (Web site). <http://www.chem.vt.edu/RVGS/ACT/notes/chem-eqm.html> (June 9, 2001).
"Chemical Sciences: Mechanism of Catalysis." University of Alberta Department of Chemistry (Web site). <http://www.chem.ualberta.ca/~plambeck/che/p102/p02174.htm> (June 9, 2001).
Ebbing, Darrell D.; R. A. D. Wentworth; and James P. Birk. Introductory Chemistry. Boston: Houghton Mifflin, 1995.
Hauser, Jill Frankel. Super Science Concoctions: 50 Mysterious Mixtures for Fabulous Fun. Charlotte, VT: Williamson Publishing, 1996.
"Mark Rosen's Chemical Equilibrium Links" (Web site). <http://users.erols.com/merosen/equilib.htm> (June 9, 2001).
Oxlade, Chris. Chemistry. Illustrated by Chris Fairclough. Austin, TX: Raintree Steck-Vaughn, 1999.
Zumdahl, Steven S. Introductory Chemistry: A Foundation, 4th ed. Boston: Houghton Mifflin, 2000.
The minimal energy required to convert reactants intoproducts, symbolized E a.
A mixture of water and a substance that is dissolved in it.
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.
A representation of a chemical reaction in which the chemical symbols on the left stand for there actants, and those on the right are the product or products.
A situation in which the ratio between the reactants and products in a chemical reaction is constant, and in which the forward reactions and reverse reactions take place at the same rate.
A process whereby the chemical properties of a substance are changed by a rearrangement of the atoms in the substance.
A generic term used for any substance studied in chemistry—whether it be an element, compound, mixture, atom, molecule, ion, and so forth.
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.
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 reformation of atoms.
A term describing a chemical reaction in which heat is absorbed or consumed.
A term describing a chemical reaction in which heat is produced.
A chemical reaction symbolized by a chemical equation in which the reactants and product are separated by an arrow that points to the right, toward the products.
Chemical equilibrium in which the substances involved are in different phases of matter (solid, liquid, gas.)
Chemical equilibrium in which the substances involved are in the same phase of matter.
LE CHÂTELIER'S PRINCIPLE:
Astatement, formulated by French chemist Henri Le Châtelier (1850-1936), which holds that whenever a stress or change is imposed on a system in chemical equilibrium, the system will adjust the amounts of the various substances in such a way as to reduce the impact of that stress.
The substance or substances that result from a chemical reaction.
A substance that interacts with another substance in a chemical reaction, resulting in a product.
A chemical reaction symbolized by a chemical equation in which the products of a forward reaction have become the reactants, and the reactants of the forward reaction are now the products. This is indicated by anarrow that points toward the left.
In chemistry and other sciences, the term "system" usually refers to any set of interactions isolated from the rest of the universe. Anything outside of the system, including all factors and forces irrelevant to a discussion of that system, is known as the environment.
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Chemical equilibrium (plural equilibria) is a dynamic condition (meaning it is marked by continuous change) in which the rate at which two opposing chemical changes is the same. As an example, consider the reaction in which ammonia gas (NH3) is made from the elements nitrogen (N2) and hydrogen (H2). That reaction can be represented by the following chemical equation:
N2 + 3 H2 ⇆ 2 NH3
The double arrow (⇆) in this equation means that two reactions are taking place at the same time. In one reaction, nitrogen and hydrogen combine to form ammonia:
N2 + 3 H2 → 2 NH3
In the second reaction, ammonia breaks down to form nitrogen and hydrogen. This reaction is just the reverse of the first reaction:
2 NH3 → N2 + 3 H2
The term chemical equilibrium refers to the condition in which both of the above two reactions are taking place at the same time.
It is easy to show why many kinds of chemical reactions must reach a point of chemical equilibrium. In the above example, suppose that the reaction begins when nitrogen gas and hydrogen gas are mixed with each other. At that moment in time, reaction number (1) takes place, but reaction number (2) is impossible. No ammonia exists at the beginning of the reaction, so equation (2) cannot occur.
Words to Know
Concentration: The amount of a substance present in a given volume, such as the number of molecules in a liter.
Dynamic condition: A condition in which components are constantly changing.
Precipitate: A solid formed during a chemical reaction.
As time goes on, the rate of reaction (1) continues to be high. A lot of nitrogen and hydrogen are available to keep the reaction going. But now reaction (2) can begin to occur. As ammonia is formed, some of it can begin to break down to form the original gases—nitrogen and hydrogen. At this point, we can say that the rate of reaction (1) is greater than the rate of reaction (2).
Over time, as nitrogen and hydrogen are used up to form ammonia, the rate of reaction (1) slows down. At the same time, the amount of ammonia gets larger and the rate of reaction (2) becomes greater. Eventually, the two rates will be equal to each other: the rate of reaction number (1) will equal the rate of reaction number (2). The system has reached a state of chemical equilibrium.
What happens if the rate of reaction (1) continues to increase beyond equilibrium? That statement means that more and more hydrogen and nitrogen are used up until they are both gone. In other words, the reaction has gone to completion. That result can occur, but it usually does not take place in chemical reactions.
Consider what happens if the rate of reaction (2) becomes greater than the rate of reaction (1). That means that ammonia breaks down faster than it is being produced. At some point, all the ammonia will be gone, and only nitrogen and hydrogen will be left. So it becomes obvious that in many chemical reactions, a point of equilibrium must be reached.
The conditions under which a chemical equilibrium exists can change, thereby changing the equilibrium itself. In general, equilibria are sensitive to three factors: temperature, pressure, and concentration. Consider once again the reaction between nitrogen and hydrogen to form ammonia:
N2 + 3 H2 ⇆ 2 NH3
What happens to this equilibrium if the temperature is increased? An increase in temperature increases the rate at which molecules move. The faster molecules move, the more likely they are to react with each other. In the above example, increasing the temperature increases the likelihood that nitrogen and ammonia molecules will react with each other and the rate of reaction number (1) will increase. The rate of reaction (2) will not change. Eventually a new equilibrium will be established reflecting this change of reaction rates.
Changing the pressure on a reaction involving gases produces a similar effect. Increasing the pressure brings molecules more closely together and increases the chances of their reacting with each other.
Finally, changing the concentration (number of molecules present) of substances in the reaction can change the equilibrium. Suppose that a lot more hydrogen is added to the previous reaction. With more hydrogen molecules present, the rate of the forward reaction will increase. Again, a new equilibrium will be reached that reflects this changed rate of reaction.
Reactions that go to completion
Most chemical reactions can be described by the previous explanation. Some cannot. Various factors can force a reaction not to reach equilibrium; instead, the reaction is said to go to completion. The phrase go to completion means that the forward direction—such as reaction (1) above—continues until all reactants are used up. The product is prevented from breaking down—as in reaction (2) above—to form the original reactants.
One condition that leads to a completed reaction is the formation of a gas that escapes from the reaction. When zinc metal (Zn) is added to hydrochloric acid (HCl), for example, hydrogen gas (H2) is formed. The hydrogen gas bubbles away out of the reaction. Since it is no longer present, the reverse reaction cannot occur:
Zn + 2 HCl → ZnCl2 + H2 ↑
(The upward-pointing arrow in the equation means that hydrogen escapes as a gas.)
Another condition that leads to a completed reaction is the formation of a precipitate in a reaction. A precipitate is a solid that forms during a chemical reaction. When silver nitrate (AgNO3) is added to hydrochloric acid (HCl), silver chloride (AgCl) is formed. Silver chloride is insoluble and settles out of the reaction as a precipitate. Since the silver chloride is no longer present in the reaction itself, the reverse reaction (AgCl + HNO3 → AgNO3 + HCl) cannot occur:
AgNO3 + HCl → AgCl ↓ + HNO3
(The downward-pointing arrow means that silver chloride forms as a precipitate in the reaction.)
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chemical equilibrium, state of balance in which two opposing reversible chemical reactions proceed at constant equal rates with no net change in the system. For example, when hydrogen gas, H2, and iodine gas, I2, are mixed, and gaseous hydrogen iodide, HI, is formed according to the equation H2 + I2 → 2HI, no matter how long the reaction is allowed to proceed some quantity of hydrogen and iodine will remain unreacted. The reason reactants in a reversible reaction are never completely converted to product is that an opposing reaction is taking place simultaneously, i.e., some of the newly formed HI is being converted back into hydrogen and iodine. For any particular temperature, a point of equilibrium is reached at which the rates of the two opposing reactions are equal and there is no further change in the system. This equilibrium point is characterized by specific relative concentrations of reactants and products and will also be reached from the opposite direction, i.e., if one starts with hydrogen iodide and allows it to decompose into hydrogen and iodine. The equilibrium point can be described by the mass action expression, which defines the equilibrium constant, Keq, in terms of the ratio of the molar concentrations of the products to those of the reactants. For the reversible reaction used as an example, the equilibrium constant is Keq=[HI]2/[H2][I2]; for the general reversible reaction nA + mB + · · · ⇌ pC + qD + · · · , the equilibrium constant is: where [A], [B], [C], [D], … are the molar concentrations of the substances and n,m, p, q, … are the coefficients of the balanced chemical equation. The larger the equilibrium constant for a given reaction, the more the reaction is favored, since a larger value of Keq means larger concentrations of the products relative to the reactants. The equilibrium constant is related to the change in the standard free energy, G°, of the system by the equation ΔG° = -RT. ln Keq, where R is a constant, T is the temperature in degrees Kelvin, and ln Keq is the natural logarithm of the equilibrium constant. Chemical equilibrium can be defined for many types of chemical processes, such as dissociation of a weak acid in solution, solubility of slightly soluble salts, and oxidation-reduction reactions. In all of these cases, the equilibrium constant or its analogue is defined for certain conditions of temperature and other factors. If any of these factors change, the system will respond to establish a new equilibrium, in accordance with Le Châtelier's principle.
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