Collisions between reactants

The dependence of reaction rate on concentration is readily explained. Ordinarily, reactions occur as the result of collisions between reactant molecules. The higher the concentration of molecules, the greater the number of collisions in unit time and hence the faster the reaction. As reactants are consumed, their concentrations drop, collisions occur less frequently, and reaction rate decreases. This explains the common observation that reaction rate drops off with time, eventually going to zero when the limiting reactant is consumed. 

There are a couple of reasons why collision between reactant molecules does not always lead to reaction. For one thing, the molecules have to be properly oriented with respect to one another when they collide. Suppose, for example, that the carbon atom of a CO molecule... 

Sonoelectrochemistry has also been used for the efficient employment of porous electrodes, such as carbon nanofiber-ceramic composites electrodes in the reduction of colloidal hydrous iron oxide [59], In this kind of systems, the electrode reactions proceed with slow rate or require several collisions between reactant and electrode surface. Mass transport to and into the porous electrode is enhanced and extremely fast at only modest ultrasound intensity. This same approach was checked in the hydrogen peroxide sonoelectrosynthesis using RVC three-dimensional electrodes [58]. 

Where R is the gas constant (8-32 joules mol-1 deg 1), and A is a constant for the reaction—independent of temperature—that is related to the proportion of the total number of collisions between reactant molecules that result in successful conversion into products. The value... 

Simple collision theory recognizes that a collision between reactants is necessary for a reaction to proceed. Does every collision result in a reaction Consider a 1 mL sample of gas at room temperature and atmospheric pressure. In the sample, about 10 collisions per second take place between gas molecules. If each collision resulted in a reaction, all gas phase reactions would be complete in about a nanosecond (10 s)—a truly explosive rate As you know from section 6.2, however, gas phase reactions can occur quite slowly. This suggests that not every collision between reactants results in a reaction. 

In order for a collision between reactants to result in a reaction, the collision must be effective. An effective collision—one that results in the formation of products—must satisfy the following two criteria. You will investigate these criteria over the next few pages. 

In this section, you used collision theory and transition state theory to explain how reaction rates are affected hy various factors. You considered simple reactions, consisting of a single-step collision between reactants. Not all reactions are simple, however. In fact, most chemical reactions take place via several steps, occurring in sequence. In the next section, you will learn about the steps that make up reactions and discover how these steps relate to reaction rates. 

State two requirements for an effective collision between reactants. 

Transition Complexes. The numerous collisions between reactant molecules result in a wide distribution of energies among the individual molecules. This can result in strained bonds, unstable forms of molecules, or unstable association of molecules which can then either decompose to give products, or by further collisions return to molecules in the normal state. Such unstable forms are called transition complexes. 

By contrast, when both the reactive solute molecules are of a size similar to or smaller than the solvent molecules, reaction cannot be described satisfactorily by Langevin, Fokker—Planck or diffusion equation analysis. Recently, theories of chemical reaction in solution have been developed by several groups. Those of Kapral and co-workers [37, 285, 286] use the kinetic theory of liquids to treat solute and solvent molecules as hard spheres, but on an equal basis (see Chap. 12). While this approach in its simplest approximation leads to an identical result to that of Smoluchowski, it is relatively straightforward to include more details of molecular motion. Furthermore, re-encounter events can be discussed very much more satisfactorily because the motion of both reactants and also the surrounding solvent is followed. An unreactive collision between reactant molecules necessarily leads to a correlation in the motion of both reactants. Even after collision with solvent molecules, some correlation of motion between reactants remains. Subsequent encounters between reactants are more or less probable than predicted by a random walk model (loss of correlation on each jump) and so reaction rates may be expected to depart from those predicted by the Smoluchowski analysis. Furthermore, such analysis based on the kinetic theory of liquids leads to both an easy incorporation of competitive effects (see Sect. 2.3 and Chap. 9, Sect. 5) and back reaction (see Sect. 3.3). Cukier et al. have found that to include hydrodynamic repulsion in a kinetic theory analysis is a much more difficult task [454]. 

Why don t all collisions between reactant molecules lead to product formation ... 

As discussed earlier, the products of the reaction are formed as a result of the collisions between reactant particles. There are more particles in a more concentrated solution and the collision rate between reactant particles is higher. The more often the particles collide, the greater the chance they have of having sufficient energy to overcome the activation energy of the reaction, and of a successful collision occurring. This means that the rate of a chemical reaction will increase if the concentration of reactants is increased, because there are more particles per unit volume. 

The events related to the suspace PaP > would correspond to direct collisions between reactants that may either leave the system as it is or it may interconvert into products via the interaction hamiltonian W. The component Q LP > contains events involving states in the orthogonal complement that includes supermolecule states. To populate these states require physical excitation processes these states may have finite lifetimes. The Q-component is called the time-delaying component by Feshbach in the context of nuclear reactions [29]. The Q T > component is given by... 

Reaction rates are characterized by the number of successful collisions between reactants, which is represented by the product of the total number of collisions of the reactants, the number of collisions that have sufficient energy to cause a reaction event (energy factor), and the fraction of collisions that have the proper orientation (probability factor). The theoretical reaction rate (k) is given by... 

The Activation Energy of Chemical Reactions Only a small fraction of the collisions between reactant molecules convert the reactants into the products of the reaction. This can be understood by turning, once again, to the reaction between C1N02 and NO. 

Theory of reaction rates that states that effective collisions between reactant molecules must occur in order for the reaction to occur. 

In order to predict the value of the frequency factor, one may assume that all collisions between reactant molecules with sufficient activation energy result in the instantaneous formation of the reaction products. With this simple hypothesis (collision theory), if the activation energy is known, then the problem of computing the reaction rate reduces to the problem of computing the rate of collision between the appropriate reactant molecules in the ideal gas mixture. This last problem is easily solved by the elementary kinetic theory of gases. 

The forward reaction is favored by a low temperature because the stress caused by the heat generated hy the reaction is reduced. But low temperature decreases the number of collisions between reactants, thus decreasing the rate of reaction. Haher compromised hy using an intermediate temperature of about 450°C. 

Increased number of collisions between reactants per second ... 

Increased fraction of collisions between reactants with the minimum energy to react... 

Explain why a collision between reactant particles must have a certain minimum energy (activation energy) in order to proceed to products. 

The reactions listed below are run at the same temperature. The activation energy for the first reaction is 132 kj/mol. The activation energy for the second reaction is 76 kj/mol. In which of these reactions would a higher fraction of collisions between reactants have the minimum energy necessary to react (the activation energy) Explain your answer. 

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