Equilibrium, chemical

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.

Explanation

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.

Upsetting equilibria

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.)

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, H ₂ , and iodine gas, I ₂ , are mixed, and gaseous hydrogen iodide, HI, is formed according to the equation H ₂  + I ₂  → 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, K_eq , 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 K_eq =[HI] ² /[H ₂ ][I ₂ ]; for the general reversible reaction n A +  m B + · · · ⇌  p C +  q D + · · · , 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 K_eq 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 Δ G ° = - RT. ln K_eq , where R is a constant, T is the temperature in degrees Kelvin, and ln K_eq 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 .

chemical equilibrium Balance in a reversible reaction, when two opposing reactions proceed at constant equal rates with no net change in the system. The initial rate of the reactions falls off as the concentrations of reactants decrease and the build-up of products causes the rate of the reverse reaction to increase. This is seen in the Haber process where ammonia formed by nitrogen and hydrogen immediately breaks down again into those elements.

chemical equilibrium see chemical equilibrium .