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Chemical kinetics



Chemical kinetics, also known as reaction kinetics, is the study of rates of chemical processes. Chemical kinetics includes investigations of how different experimental conditions can influence the speed of a chemical reaction and yield information about the reaction's mechanism and transition states. In 1864, Peter Waage and Cato Guldberg pioneered the development of chemical kinetics by formulating the law of mass action, which states that the speed of a chemical reaction is proportional to the quantity of the reacting substances.  

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Contents

Rate of reaction

Main article: reaction rate

Chemical kinetics deals with the experimental determination of reaction rates from which rate laws and rate constants are derived. Relatively simple rate laws exist for zero order reactions (for which reaction rates are independent of concentration), first order reactions, and second order reactions, and can be derived for others. In consecutive reactions the rate-determining step often determines the kinetics. In consecutive first order reactions, a steady state approximation can simplify the rate law. The activation energy for a reaction is experimentally determined through the Arrhenius equation and the Eyring equation. The main factors that influence the reaction rate include: the physical state of the reactants, the concentrations of the reactants, the temperature at which the reaction occurs, and whether or not any catalysts are present in the reaction.

Factors affecting reaction rate

Nature of the Reactants

Depending upon what substances are reacting, the time varies. Acid reactions, the formation of salts, and ion exchange are fast reactions.When covalent bond formation takes place between the molecules and when large molecules are formed, the reactions tend to be very slow.

Physical State

The physical state (solid, liquid, or gas) of a reactant is also an important factor of the rate of change. When reactants are in the same phase, as in aqueous solution, thermal motion brings them into contact. However, when they are in different phases, the reaction is limited to the interface between the reactants. Reaction can only occur at their area of contact, in the case of a liquid and a gas, at the surface of the liquid. Vigorous shaking and stirring may be needed to bring the reaction to completion. This means that the more finely divided a solid or liquid reactant, the greater its surface area per unit volume, and the more contact it makes with the other reactant, thus the faster the reaction. To make an analogy, for example, when you start a fire, first you use wood chips and small branches - you don't start with big logs right away. In organic chemistry On water reactions are the exception to the rule that homogeneous reactions take place faster than heterogeneous reactions.

Concentration

Concentration plays an important role in reactions. According to the collision theory of chemical reactions, this is because molecules must collide in order to react together. As the concentration of the reactants increases, the frequency of the molecules colliding increases, striking each other faster by being in closer contact at any given point in time. Imagine two reactants being in a closed container. All the molecules contained within are colliding constantly. By increasing the amount of one or more of the reactants you cause these collisions to happen more often, increasing the reaction rate (Figure 1.1)..

Temperature

Temperature usually has a major effect on the speed of a reaction. Molecules at a higher temperature have more thermal energy. When reactants (reactant + reactant → product) in a chemical reaction are heated, the more energetic atoms or molecules have a greater probability to collide with one another. Thus, more collisions occur at a higher temperature, making a product in a chemical reaction. More importantly however, is the fact that at higher temperatures molecules have more vibrational energy, that is, atoms are vibrating much more violently, so raising the temperature not only increases the number of collisions but also collisions that can result in rearrangement of atoms within the reactant molecules. For example, a refrigerator slows down the speed of the rate of reaction since it cools the molecules. On the other hand, an oven gives heat (energy) to the molecules which in turn speeds up the rate of reaction, cooking the food faster.

A reaction's kinetics can also be studied with a temperature jump approach. This involves using a sharp rise in temperature and observing the relaxation rate of an equilibrium process.

Catalysts

A catalyst is a substance that accelerates the rate of a chemical reaction but remains chemically unchanged afterwards. The catalyst increases rate reaction by providing a different reaction mechanism to occur with a lower activation energy. In autocatalysis a reaction product is itself a catalyst for that reaction leading to positive feedback. Proteins that act as catalysts in biochemical reactions are called enzymes. Michaelis-Menten kinetics describe the rate of enzyme mediated reactions.

In certain organic molecules specific substituents can have an influence on reaction rate in neighbouring group participation.

Agitating or mixing a solution will also accelerate the rate of a chemical reaction, as this gives the particles greater kinetic energy, increasing the number of collisions between reactants and therefore the possibility of successful collisions.

Increasing the pressure in a gaseous reaction will increase the number of collisions between reactants, increasing the rate of reaction. This is because the activity of a gas is directly proportional to the partial pressure of the gas. This is similar to the effect of increasing the concentration of a solution.

Equilibria

While chemical kinetics is concerned with the rate of a chemical reaction, thermodynamics determines the extent to which reactions occur. In a reversible reaction, chemical equilibrium is reached when the rates of the forward and reverse reactions are equal and the concentrations of the reactants and products no longer change. This is demonstrated by, for example, the Haber-Bosch process for combining nitrogen and hydrogen to produce ammonia. Chemical clock reactions such as the Belousov-Zhabotinsky reaction demonstrate that component concentrations can oscillate for a long time before finally reaching equilibrium.

Free energy

In general terms, the free energy change (ΔG) of a reaction determines if a chemical change will take place, but kinetics describes how fast the reaction is. A reaction can be very exothermic but will not happen in practice if the reaction is too slow. If a reactant can produce two different products, the thermodynamically most stable one will generally form except in special circumstances when the reaction is said to be under kinetic reaction control. The Curtin-Hammett principle applies when determining the product ratio for two reactants interconverting rapidly, each going to a different product. It is possible to make predictions about reaction rate constants for a reaction from Free-energy relationships.

The kinetic isotope effect is the difference in the rate of a chemical reaction when an atom in one of the reactants is replaced by one of its isotopes.

Chemical kinetics provides information on residence time and heat transfer in a chemical reactor in chemical engineering and the molar mass distribution in polymer chemistry.

Applications

Reaction kinetics provides chemists with the tools to determine and describe the factors important in food decomposition, the hardening of dental materials, the reproduction of micro-organisms, the speed at which stratospheric ozone is destroyed, and how the enzymes influence chemical processes in biological systems.

See also

References

  • Preparing for the Chemistry AP Exam. Upper Saddle River, New Jersey: Pearson Education, 2004. 131-134. ISBN 0-536-73157-8
 
This article is licensed under the GNU Free Documentation License. It uses material from the Wikipedia article "Chemical_kinetics". A list of authors is available in Wikipedia.
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