Chemical reaction behavior

Type of physical science: Chemistry

Field of study: Chemical reactions

Chemical reactions are very important because humans depend on the energy from chemical reactions for countless tasks. Chemical reactions are rearrangements of atoms to produce molecules of new substances.

Overview

A chemical reaction is a process by which substances known as products are formed from substances called reactants. The atmospheric gases hydrogen and oxygen react to produce water. Sodium, a silvery metal, and the poisonous gas chlorine produce table salt. The process frequently produces substances with properties so different from the reacting materials that there is some visible sign that the change has taken place: a color change, the formation of a precipitate from solution, or the evolution of a gas. In other cases, the process is not accompanied by obvious macroscopic changes but must be monitored by sensitive instruments.

A chemical equation is a symbolic representation of the change that takes place in a reaction. The water reaction would be written:

Multiple line equation(s) cannot be represented in ASCII text; please see PDF of this article if available.

The equation gives much information about the process that takes place. The reactants and products are named by chemical formulas that tell the type and number (in subscript form) of atoms that constitute a molecule of a substance. The state (solid, liquid, or gas) of each substance is indicated in parentheses after its name. Some reactions occur much more rapidly in the presence of a catalyst. Since this substance is neither a reactant nor a product, its symbol is frequently placed over the reaction arrow (Pt = platinum in the equation above). Chemical reactions are rearrangements of atoms to produce molecules of new substances. All atoms that appear on the reactant side of the process must also appear on the product side. This principle governs the ratio of the number of molecules of each substance that react. The stoichiometric coefficients placed before the symbols for the reacting species and products indicate this ratio.

In many cases, the idea that reactants are turned into products is an oversimplification.

An important industrial example is the synthesis of ammonia from nitrogen and hydrogen (the Haber process). If hydrogen and nitrogen are placed in a reaction vessel under the conditions of high pressure and temperature, and with a catalyst, ammonia will gradually be formed. The production of ammonia ceases long before the reactants are used up. In many reactions such as this, the percentage of reactants converted to products is very small. In the water reaction, the amount of unreacted hydrogen and oxygen is minuscule. No reaction actually goes to completion, rather, it reaches an equilibrium state, in which the reaction mixture contains reactant and product molecules. If the reaction began with ammonia in the vessel, under the Haber conditions, the equilibrium mixture would contain the same ratio of ammonia to hydrogen and oxygen as it would starting from the other end. The equilibrium condition is specified by the use of a double arrow in the chemical equation: N₂(g) + 3H₂(g) ↔ 2NH₃(g)

The relative amounts of reactants and products present in a reaction system at equilibrium are given by an equilibrium constant, which is fixed for a given set of equilibrium conditions (temperature, pressure, and so on). For the ammonia reaction, the equilibrium constant would be expressed as:

Multiple line equation(s) cannot be represented in ASCII text; please see PDF of this article if available.

The brackets in the expression indicate the equilibrium concentrations of the various components of the reaction mixture. Since this particular example has species that are all gases, the equilibrium constant may be expressed in terms of the pressure exerted by each gas. The exponent of the pressure, or concentration term, is the stoichiometric coefficient of the balanced equation.

So far, the reactions mentioned have been assumed to occur in closed systems. No product or reactant could leave or be added to the system. The only reaction that would occur would be that which brings the system to the equilibrium state. If some method could be devised that would allow the product (for example, ammonia) to be gradually removed as it was formed, then the reaction would keep moving in the forward direction in an effort to reach the equilibrium state, and might go to completion.

The equilibrium state in a closed system is a dynamic rather than a static condition.

Although there are no changes in the concentrations of the various species in the reaction mixture, once the state of equilibrium is reached, both forward and reverse reactions are actually occurring, but at the same speed. The behavior of a system at equilibrium is described by Le Chatelier's principle, which states that a system will respond to changes by moving to a new equilibrium position that gives the same value of the equilibrium constant. If hydrogen or nitrogen were added to the ammonia system, the reaction would proceed forward to use the added material. Pressure or temperature variations also affect the position of equilibrium, since they affect the value of the equilibrium constant. A pressure increase (volume decrease) favors the production of ammonia, since in the balanced equation the product side has fewer molecules and therefore requires less space. Temperature increase favors the reactant side since the forward reaction produces heat as a by-product. Generally, some compromise must be reached between pressure and temperature to maximize the production of the desired species in an industrial process.

The question of why a system moves toward a certain equilibrium state is an interesting but complex one, and is the question addressed by the field of thermodynamics. Two important and general tendencies are followed by chemical reaction systems: the movement toward minimum energy, and toward maximum randomness. Most chemical and physical processes either require energy in the form of heat (are endothermic) or release energy (are exothermic).

For a reaction carried out under the condition of constant pressure, the heat of the reaction is known as the enthalpy change for the process (Δ H). At room temperature, exothermic (Δ H less than 0) reactions are generally spontaneous (occur in a system left to itself). The enthalpy criterion is not enough to determine the direction of spontaneous change. The melting of ice (an endothermic process, Δ H greater than 0) occurs spontaneously at room temperature. To grasp the concept of entropy (randomness), consider the melting of ice. In the solid state, the water molecules are held in place in a crystalline structure, with little freedom to do any moving except small amplitude vibrations in their assigned positions in the lattice. Water molecules in the liquid state are held relatively close together by strong intermolecular forces, but are able to move past one another rather freely. This additional freedom to move about results in an increase in the entropy of the system. Those processes that result in entropy increases are spontaneous. An obvious extension of this idea is to consider the freedom of motion of the same molecules once they pass into the gaseous state. The evaporation process is endothermic but results in greater entropy, which is the deciding factor in the spontaneous evaporation of water at room temperature.

The two tendencies (minimum energy and maximum entropy) frequently work against each other in determining the direction that a given reaction will take under certain conditions of temperature and pressure. A combination of entropy and enthalpy that has been found useful in predicting the direction of spontaneous change is the Gibbs energy function: δG = δH – TδS (T = temperature)

Whenever the Gibbs energy is negative, the reaction will be spontaneous. The temperature factor in the second term on the right of the above equation allows entropy to be the determining factor at high temperatures, while enthalpy is more important at low temperatures.

This explains why most reactions that occur spontaneously at room temperature are exothermic.

An important relationship exists between the value of the Gibbs energy for a reaction and its equilibrium constant: δG = -2.303RTlogK_eq (R = a constant)

Reactions having a very large negative value for the Gibbs energy will have equilibrium constants much greater than one (products strongly favored) and will essentially go to completion. An example of a reaction that has a very large equilibrium constant, about 10 to the power of 60, is that between the toxic gases carbon monoxide and nitrogen oxide to produce carbon dioxide and nitrogen. Since this reaction is predicted by its equilibrium constant to be spontaneous, it should provide a way to remove these toxic products of automobile combustion from the atmosphere. Unfortunately, there is a practical difficulty, in that the reaction occurs at an extremely slow rate. This is a quite common situation; many reactions that should in principle go to completion often take place very slowly.

There is, in fact, no correlation between the equilibrium constant and the rate of a chemical reaction. The principles of chemical kinetics help predict how rapidly a chemical reaction will occur. The rate of any chemical reaction has been found to depend in a complex way on the concentrations of the various species present in the reaction mixture. Each rate expression contains a rate constant, k, which is specific to the reaction. A knowledge of this dependence not only allows the actual rate to be predicted but also gives clues to the step-by-step mechanism by which the reaction occurs. The study of reaction rates with and without a catalyst present has aided in determining the mechanism by which the catalyst functions in increasing or decreasing the rate of the reaction. One of the most important examples is the functioning of the biological catalysts, the enzymes, without which most reactions of biological importance would occur extremely slowly.

Applications

Chemistry is a major contributor to the technological world humans are in the process of creating. Understanding of reaction processes and the ability to control them have resulted in a multitude of new materials. Drugs, plastics, ceramics, detergents, synthetic textiles, and rubber are only a few.

Knowledge of chemical kinetics has led to important developments in the understanding of the mechanism of some diseases. Enzymes in living systems are substances that catalyze almost every reaction of importance to the living organism. Unfortunately, not all enzymes work for the good of the organism. Some diseases are in fact caused by enzymes that catalyze undesirable reactions. Enzymes function by forming a complex with a molecule known as its substrate. Many enzymes will form this complex with only one type of substrate molecule, and are therefore extremely specific in their activity. Substrate molecules for an enzyme of interest have been synthesized that differ from the normal substrate in that a heavier isotope is substituted for one of the atoms normally found in the molecule. The new molecule behaves chemically like the normal molecule, except that the rate of the enzyme-catalyzed reaction is frequently changed by the substitution of the heavier isotope. Study of the changes in reaction rate as isotopic substitutions are made in various parts of the substrate molecule have provided clues to the chemical structure of the substrate while it is in the complex with the enzyme. This information, in some cases, allows the design of an alternate substrate with similar structure that might potentially be able to form a stronger complex with the enzyme. This effectively prohibits the enzyme from performing its original, undesirable activity, and is a possible method of disease control for those diseases whose symptoms are enzyme-related. This is one of many possible ways that drug therapy is used in disease control.

One of the more common elements on earth is nitrogen, which constitutes almost 80 percent of the atmosphere. In its atmospheric form, nitrogen is very stable and unreactive, but can be converted by bacteria found in some plant roots to building blocks for more complex chemical substances, such as proteins. This process is called nitrogen fixation. Before it can be used industrially, atmospheric nitrogen must be converted to a form that reacts more easily, such as ammonia. Ammonia is the starting material for such diverse products as fertilizers and explosives. Early in the 1900's, the German chemist Fritz Haber developed a process for the production of ammonia from hydrogen and atmospheric nitrogen. The ammonia reaction is one that has an equilibrium position that favors the reactants rather than the desired product, ammonia. Haber studied the effects of temperature and pressure on the position of the ammonia equilibrium, and maximized its production. This accomplishment allowed the German government to wage war without fear of losing the supply of nitrates normally imported from Chile, since Haber's work made Germany self-sufficient in the production of this important compound.

Humans depend on the energy from chemical reactions to perform countless tasks. All but a small fraction of the energy people use comes from the chemical combustion of fossil fuels such as coal oil or natural gas. The operation of cars, the heating or cooling of homes, and the manufacture of all types of products use fossil fuel energy. Many products--cameras, radios, hearing aids--operate using the electric power generated by the chemical reactions taking place inside batteries. Demand for continued supplies of energy-producing materials is placing a severe stress on the world's resources and is a battleground between the forces of development and conservation. The task of chemical thermodynamics is to develop new energy sources and to improve the efficiency of those processes already in use. Another important task of thermodynamics is to predict the direction that a chemical system will take in reaching equilibrium. Calculations of this type must be performed for each new proposed chemical process to decide its feasibility before even small-scale testing in a laboratory begins. Theoretical work using information from thermodynamics will prevent many work hours from being wasted and the huge loss of capital frequently involved in a new chemical process venture.

Context

It is impossible to separate the study of chemical reaction behavior from the development of the science of chemistry. The early Egyptians developed glues for making papyrus and plastic cement to seal coffins. Metallurgists developed techniques for separating metals from their ores and for combining some metals to produce alloys. Alchemists discovered interesting and useful techniques and properties of materials in their unsuccessful search for ways of changing nonprecious metals into gold. Out of these ancient accomplishments, partly religious, partly magic, chemistry and the chemical arts developed. Chemical theory has a later birth.

The understanding of the atomic nature of matter and the way that atoms combined and rearranged themselves to produce new and different compounds began with the discovery of oxygen by Antoine Lavoisier and others. John Dalton carefully reproduced the experiments of Lavoisier and made many careful measurements of the masses of materials that combined to make various compounds. Dalton found that compounds were formed from elements that always combined in a fixed weight ratio. Water, by weight, was eight parts oxygen to one part hydrogen.

He used HO as the formula for water, and concluded that each oxygen atom must be eight times the mass of each hydrogen atom. The numerical conclusions were incorrect, but the thinking was in the right direction, and the development of chemical theory began.

Physical chemistry (the home of thermodynamics, kinetics, and equilibrium studies) developed at the beginning of the nineteenth century. At this time, physicists such as James Joule, Julius Robert von Mayer, and Hermann von Helmholtz were studying the flow of energy that is heat. Nicolas-Leonard-Sadi Carnot, theoretically studying heat engines, and Sir William Thomson (Lord Kelvin) and Rudolf Clausius showed that heat flowed spontaneously along a temperature gradient. These developments in physics were not isolated from chemistry, since the important sources of heat in the nineteenth-century world were in chemical reactions such as fuel combustion. The fields of chemistry and physics met around 1840 as a result of the research of Germain Henri Hess, who measured the heat resulting from a number of chemical reactions.

Investigation into the spontaneity of reactions followed shortly afterward, when Marcelin Berthelot devised a water calorimeter for measuring reaction heats. He ran careful determinations of the heats of hundreds of reactions, and proposed that spontaneous reactions were those that gave off heat.

The history of the study of chemical reactions is long and varied, and the field has a future as full of promise as its past is full of accomplishment. Research in the field of chemical reaction behavior will continue in the search for new and better fuels, drugs, and synthetic substances of many types, from fabrics to structural materials. In some sense, chemistry has lost none of its magic.

Bibliography

Asimov, Isaac. A SHORT HISTORY OF CHEMISTRY. Garden City, N.Y.: Doubleday, 1965. Asimov's easy style has made him one of the most successful of all science writers. In this book, he manages to present the important advances in the field of chemistry in an entertaining and nontechnical manner.

Fenn, John B. ENGINES, ENERGY, AND ENTROPY. San Francisco: W. H. Freeman, 1982. This is not an elementary text in the sense that some mathematical background beyond algebra would be helpful. It is a clear if sometimes whimsical exposition of the rather abstract topic of thermodynamics.

Irwin, Keith Gordon. THE ROMANCE OF CHEMISTRY. New York: Viking Press, 1959. Although dated, it is a worthwhile book. It touches in a brief way the important events in the history of chemistry from alchemy to nuclear fission. The important part played by the study of reaction behavior in the development of chemistry as a science is clear.

Selinger, Ben. CHEMISTRY IN THE MARKETPLACE. San Diego: Harcourt Brace Jovanovich, 1988. This is a reference book for the general reader.

Selinger deals with all types of consumer chemistry and the relationship between chemistry and society. The appendix on the entropy game is excellent for presenting the concept of entropy.

Turner, A. Mason, and Curtis T. Sears, Jr. INQUIRIES IN CHEMISTRY. Boston: Allyn & Bacon, 1974. This is an interesting and understandable high school level chemistry textbook. The chapters on energy, chemical reaction rates, and equilibrium are easy to read and fairly nonmathematical. There are even some simple experiments using common household materials for the more adventurous.

Principal terms

Quantum Mechanics of Chemical Bonding

Chemical Formulas and Combinations

Chemical Reactions and Collisions

ATOM: the smallest particle of matter that characterizes an element

EQUILIBRIUM: a condition in which forward and reverse processes proceed at equal rates and no further net change occurs

ISOTOPES: forms of an element that differ only in the mass of the atom (for example, the atomic weight of chlorine is either 35 or 37)

KINETICS: the measurement, control, and prediction of the rate of a chemical reaction

MECHANISM: a series of steps by which reactants are converted into products in a chemical reaction

MOLECULE: a combination of atoms that can exist as an individual identifiable unit possessing a unique set of measurable properties

PRECIPITATE: a solid that separates from a liquid solution

STOICHIOMETRY: measurements and relationships involving substances and mixtures of chemical interest

THERMODYNAMICS: the study of the relationship between chemical reactions and energy