Chemical reactions and collisions

Type of physical science: Chemistry

Field of study: Chemistry of molecules: nature of chemical bonds

Almost all chemical reactions take place between molecules that, at the moment of reaction, have energies far greater than the average energies of the molecules around them. This energy usually accumulates in a given molecule as a result of its collisions with other molecules. Energy requirements of a given reaction are usually described in terms of a potential energy surface with a high point over which the reacting molecule(s) must pass.

Overview

All atoms are in constant motion. This is true at all temperatures and does not depend on whether the atoms are in a solid, liquid, or gas. It is just as true for atoms in molecules as for atoms moving alone. As the temperature increases, the average value of the energy of motion--kinetic energy--increases. This means that the atoms and molecules move faster as the temperature rises. As an illustration of such motion, the most likely speed of a nitrogen molecule, consisting of two nitrogen atoms, is slightly more than 400 meters per second in air at room temperature. There are three reasons why there is no sense of such rapid motion. First, each molecule is extremely light, so that its impact on humans cannot be felt. Second, there are vast numbers of molecules moving randomly in all directions around people. Third, molecules frequently collide with one another and change directions. The typical nitrogen molecule will undergo roughly a billion collisions per second in air.

Molecules exchange energy through collisions. At a given moment, only a small fraction of the molecules will have exactly the kinetic energy required for speeds between 400 and 401 meters per second. The next moment, other molecules will have the same amount of energy. As long as the temperature is constant, the kinetic energy of a typical molecule will remain the same. Yet, individual molecules are continually exchanging kinetic energy through collisions. Some of them will have more than the typical molecule, and some less. The number of molecules at some high energy states decreases as that energy increases. Normally, only the molecules with kinetic energies far above average are able to react.

Chemical bonds are strong enough that a normal molecule is not damaged by collisions at ordinary temperatures. No reactions occur between the nitrogen and oxygen molecules of the air, despite their ceaseless banging together and bouncing apart. Also, neither the nitrogen molecules nor the oxygen molecules react with the water vapor in the air.

A chemical reaction requires enough energy to weaken bonds, but this does not happen in the typical collision. Instead, it requires a much more energetic collision. Consider the case of two molecules that are about to collide. If they possess energy far in excess of the average as a result of immediately preceding collisions, their collision with each other can be much more energetic than the typical one. If it is, a bond can be weakened, and a reaction may take place.

In order to have a simple sketch of a reaction, let A-B be a molecule, and C an atom. In this hypothetical reaction, the atom B is to leave A and attach to C: A-B + C → A + B-C

One could imagine the reaction to occur in a simple manner: The molecule A-B gains enough energy through collision to break apart, after which atoms B and C come together to make B-C. If that were the case, the collisions would have to provide all the energy needed to break the bond between A and B. A schematic potential energy surface for such a reaction would look like this:

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The difference between the reaction in progress and the reactants before the reaction is that A-B has been broken into separate A and B atoms. This requires much energy, and it must come from collisions.

This simple description is inaccurate for most reactions because it overlooks an important factor: C must have some attraction for B. Otherwise C and B would not join to make the product molecule B-C. Atom C will already have some attraction to B while B is attached to A. That attraction will not be strong enough to pull B away from A unless the molecule A-B is weakened by collision. With atom C pulling B toward itself, a more realistic potential energy surface is as follows:

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This potential energy surface differs in two respects from the previous one. First, the reaction in progress is represented by a weakened, longer bond between A and B, that is, A--B, while the bond between B and C is forming before A--B is completely broken. Second, this simultaneous bond breaking and bond making lowers the energy that must be supplied in order to move up to the reaction in progress.

This reduction of the energy required for reaction has great consequences for chemical reactions. Since molecules pass energy back and forth by collision, and since few molecules have energies far in excess of the average at a given moment, the result of taking into account the bond formation between C and B, while the bond between A and B is breaking, is to increase the small number of molecules that have the energy needed for reaction.

The energy needed to move from reactants before reaction up to the reaction in progress is called the activation energy. The lower the activation energy, the faster the reaction.

Applications

The activation energy is normally a property of the atoms and molecules that enter into reaction. In the usual case, it is far more energy than the average among molecules so that only a tiny fraction of collisions leads to reaction. One of the ways to change the activation energy is to increase the temperature. This has the effect of increasing the average energy of the atoms and molecules before reaction, thereby reducing the energy difference between inactive reactants and the reaction in progress. An increase in temperature of, for example, 40 degrees Celsius would be invisibly small on the potential energy diagram, but it would have a great influence on the rate of reaction. This is true because raising the temperature increases the tiny fraction of molecules with enough collisional energy to react faster than it raises the average energy of all molecules. This fact is one of the great principles of both physics and chemistry. It is evident in an observation that everyone has made: Water evaporates much faster after rain in the summer than in the winter. Someone may object that this has nothing to do with chemical reactions, because water on the sidewalk and water vapor in the air are the same chemical substance. Nevertheless, the rate of any change that depends strongly on the temperature obeys the same law of nature, whether it is a physical change or a chemical reaction.

One can find the value of the activation energy by studying the rate of reaction at various temperatures. Its numerical value is a property of the particular reaction. A typical activation energy is more than ten times the average kinetic energy of the molecules at the temperature of the reaction. The series of experiments needed to find the activation energy will also show that it is essentially constant for many reactions even if the rate is studied at a temperature range as great as 100 degrees Celsius. If the activation energy is more than ten times as great as the average kinetic energy, then an increase in the latter will have a minor effect on the difference between them.

The potential energy surface can be more complicated. This happens when a reaction is not completed in one step, but first makes a molecule that reacts further to make stable products.

The following reaction between gases illustrates this. The overall reaction is NO₂ + CO → NO + CO₂, but a careful study of the rate of the reaction shows that it happens in two steps: NO₂ + NO₂ → NO₃ + NO followed by NO₃ + CO → NO₂ + CO₂. The potential energy surface must have the following general appearance:

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At point 1, the reactants have their typical energies at the temperature of reaction. If they are to reach point 2w, where the first reaction is in progress, collisions must supply the activation energy. At point 2, one oxygen (O) atom passes from one molecule to the other to make a molecule of NO, a final product. Yet, the NO₃ is a reactive intermediate that collides with a molecule of CO in reaching point 4, where the second reaction is in progress. That reaction produces the other final product, CO₂, and also remakes one reactant molecule.

Each of the two reactions is considered an elementary reaction and each has its own activation energy; for the first reaction, it is the energy needed to get from point 1 to point 2. For the second reaction, it is the energy needed to get from point 3 to point 4. The diagram makes it clear that the second activation energy is far less than the first one. That the first activation energy should be greater than the second is the case for this pair of reactions, but it is not necessarily the case for all overall reactions that happen in two steps.

It is obvious that the overall reaction could have been carried out by the simple transfer of one oxygen atom from NO₂ to CO. What advantage can possibly cause nature to run this reaction in two steps? The activation energy required to go from point 1 to point 2 must be less than the activation energy required to transfer an O atom from NO₂ to CO directly in one step.

The reason for this is that excess energy is hard to acquire, and a reaction, even a relatively complicated one, that requires a lower activation energy will always be favored over one with a higher activation energy. Some of the reactions that take place in living systems are highly complicated.

It has been assumed thus far that reactions take place between molecules in gases. The collisions of interest are among the reactants. In the case of reactions in solution, the reactant is surrounded by molecules of the solvent. If it is to react, it must come into contact with another reactant molecule. Most of its collisions, however, are with the solvent. When two reactant molecules meet, they may collide more than once with each other, because the solvent molecules keep them from moving apart immediately after the initial collision.

The rate of a chemical reaction often has serious consequences, and great efforts have been made in research to alter reaction rates. Sometimes it is desirable to slow down a rate, as when a drug that improves medical care has harmful side effects. More often, however, the scientist wants to increase the speed of a reaction. An increase in temperature will in some cases bring about the added reaction rate. At times, the temperature cannot be increased, as for example if the solvent will boil or a reactant will break down at the higher temperature. Also, heating is costly, especially in a manufacturing process that is carried out on great quantities of material. In such cases, the scientist will search for a catalyst--a substance that is present in the reaction mixture and that is able to lower the activation energy without being used up. In general, a catalyst must bring the reactants together in a way that facilitates the breaking and making of bonds. If the catalyst can do this, the rate of reaction will increase, and the gain is sometimes spectacular, with great savings in production costs. Various important industrial catalysts are metals that have been prepared in a manner that gives them large surface areas.

In the realm of living systems, enzymes are the catalysts, and they often bring about tremendous rate increases for highly specific reactions. Life would be impossible without them.

Enzymes are typically large molecules and act by attaching a reactant molecule called a substrate to their surface in a way that makes the substrate enter more easily into reaction.

Context

Thus far it has been assumed that heat is the only source of energy for a reaction. There are others. If molecules are put into solution alongside a piece of metal that carries an electric charge, they can lose electrons to a positively charged metal or take them from a negatively charged one. In such a case, the reaction is being driven electrically, and the process is called electrolysis. The temperature will have little effect on it, except that it will influence the rate at which molecules move through solution. Highly stable molecules such as water can be broken apart by electrolysis. Temperatures of several thousands of degrees would be needed to break water apart by heat alone.

If a particle of light collides with a molecule, it is absorbed and adds its energy to the molecule. In many cases, particularly for ultraviolet light and visible light on the side of the spectrum close to the ultraviolet region, this added energy serves as an activation energy (though not usually described as such) that permits the molecule to react at once. The study of reactions caused by light is called photochemistry. Photographic films capture an image through the use of photochemical reactions. The fact that such films are almost unaffected by wide changes in temperature is proof that the reactions are driven by light rather than heat.

Lasers are used as a source of energy for many kinds of reactions. In an industrial reaction, for example, laser light of the right wavelength can cause a reaction to produce more of the desired product and less of the unwanted side products.

If light of sufficient energy is absorbed, it can ionize the molecule, which means that the molecule loses an electron and becomes electrically charged. This ionization is more than enough energy to activate a molecule, so that it may react on the next collision. These reactions between gaseous ions and neutral molecules take place in the upper atmosphere, where molecules receive damaging radiation from the sun, and in lightning discharges. They are also studied extensively in the laboratory. Because the loss of an electron deposits much energy in a molecule, reactions between gaseous ions and molecules frequently happen at rates that are almost independent of temperature. Because there is no activation energy, and because any electrically charged particle will be attracted to a neutral molecule, the potential energy surface moves down when the molecule and ion first come into contact. Letting I+ be the ion (electrically charged molecule) and M the neutral molecule, the potential energy diagram has the following form for a simple reaction:

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If a situation exists in which a substantial fraction of the atoms or molecules are ionized, matter under these conditions is called a plasma. While plasmas are rare on Earth, they are probably the most common form of matter in the universe, because of the great amount of energy available in stars.

Principal terms

ACTIVATION ENERGY: the energy, in excess of the average energy of reactant molecules, that is required in order for molecules to react

ATOM: the smallest and simplest unit of an element that can exist

COVALENT BOND: the sharing of electrons by two atoms, with the result that the atoms are joined together in a molecule

ELECTRON: the smallest unit of negative electrical charge

ELEMENTARY REACTION: a reaction that leads to products as a result of a single reactive collision between reactants

ENERGY: the capacity to do work; here, the important kinds of energy are collisional energy and the capacity to stretch and break covalent bonds

KINETIC ENERGY: the energy that an object has when it is in motion; it is always positive, and the kinetic energy that goes into collisions is central to the arguments that follow

POTENTIAL ENERGY: the energy that an object has because it is in a certain position; an object at the top of a hill has potential energy, which is shown by the increase in speed as it moves downhill

REACTION: the process in which one or more molecules, called reactants, gain or lose atoms to become reaction products

SOLVENT: a liquid in which substances are dissolved in a solution

Bibliography

Atkins, Peter W. GENERAL CHEMISTRY. New York: W. H. Freeman, 1989. Chapter 12 describes reaction rates. Elementary reactions and the rate-determining step in a reaction sequence are described with particular care. The collisional point of view is discussed toward the end of the chapter. Figure 12.23 illustrates beautifully the discussion in this article.

Bailar, John C., Jr., et al. CHEMISTRY. San Diego: Harcourt Brace Jovanovich, 1989. Chapter 18 presents reaction rates with some interesting examples of the steps in reactions.

Ebbing, Darrell D. GENERAL CHEMISTRY. 3d ed. Boston: Houghton Mifflin, 1990. Chapter 14 provides a careful discussion of collision theory, effects of temperature, elementary reactions, multistep reactions, and catalysis.

Gillespie, Ronald, et al. CHEMISTRY. 2d ed. Boston: Allyn & Bacon, 1989. Chapter 19 presents a clear account of the rates of chemical reactions. The figures on pages 874-875 show the change in potential energy for reactions with positive and negative energy changes.

Kotz, John C., and Keith F. Purcell. CHEMISTRY AND CHEMICAL REACTIVITY. Philadelphia: Saunders College Publishing, 1987. Chapter 22 contains a good general discussion of reaction rates. Some ball-and-stick sketches of reactions are particularly helpful.

McQuarrie, Donald A., and Peter A. Rock. GENERAL CHEMISTRY. 3d ed. New York: W. H. Freeman, 1991. Chapter 16 contains a thorough introduction to reaction rate studies. The catalytic activity of platinum is effectively presented. Chapter 7 gives an excellent presentation of gases and kinetic energy.

Dynamics of Chemical Reactions

Isotopic Effects in Chemical Reactions

Photon Interactions with Molecules