Chapter+14,+Section+3

= = S11.A.1.3.2  Describe or interpret dynamic changes to stable systems (e.g.,chemical reactions, human body, food webs, tectonics, homeostasis).

S11.A.1.3.1 Use appropriate quantitative data to describe or interpret change in systems (e.g., biological indices, electrical circuit data, automobile diagnostic systems data).

=Chemical Kinetics=

Concepts
In many reactions, the rate of reaction changes as the reaction progresses. Initially the rate of reaction is relatively large, while at very long times the rate of reaction decreases to zero (at which point the reaction is complete). In order to characterize the kinetic behavior of a reaction, it is desirable to determine how the rate of reaction varies as the reaction progresses. A **rate law** is a mathematical equation that describes the progress of the reaction. In general, rate laws must be determined experimentally. Unless a reaction is an elementary reaction, it is not possible to predict the rate law from the overall chemical equation. There are two forms of a rate law for chemical kinetics: the **differential rate law** and the **integrated rate law**. The differential rate law relates the rate of reaction to the concentrations of the various species in the system. Differential rate laws can take on many different forms, especially for complicated chemical reactions. However, most chemical reactions obey one of three differential rate laws. Each rate law contains a constant, //k//, called the **rate constant**. The units for the rate constant depend upon the rate law, because the rate always has units of mole L-1 sec-1 and the concentration always has units of mole L-1.
 * Zero-Order Reaction**

For a zero-order reaction, the rate of reaction is a constant. When the limiting reactant is completely consumed, the reaction abrupts stops.

Differential Rate Law: //r// = //k//

The rate constant, //k//, has units of mole L-1 sec-1.
 * First-Order Reaction**

For a first-order reaction, the rate of reaction is directly proportional to the concentration of one of the reactants.

Differential Rate Law: //r// = //k// [A]

The rate constant, //k//, has units of sec-1.
 * Second-Order Reaction**

For a second-order reaction, the rate of reaction is directly proportional to the square of the concentration of one of the reactants.

Differential Rate Law: //r// = //k// [A]2

The rate constant, //k//, has units of L mole-1 sec-1. These three behaviors are illustrated in the following plots. The graph at the left shows concentration-time plots for zero-order, first-order, and second-order reactions. The corresponding rate-concentration plots are shown at the right. In examining the plots, bear in mind that as the reaction progresses, the concentration of reactant decreases. This corresponds to moving from right to left on the plot of reaction rate vs concentration. In this example, the reactant has a stoichiometric coefficient of one, so the reaction rate (plotted in the graph at the right) corresponds with the negative value of the slope of the concentration-time curve (plotted in the graph at the left). Carefully examine the graphs and take note of the following points:


 * For a zero-order reaction, the rate of reaction is constant as the reaction progresses.
 * For a first-order reaction, the rate of reaction is directly proportional to the concentration. As the reactant is consumed during the reaction, the concentration drops and so does the rate of reaction.
 * For a second-order reaction, the rate of reaction increases with the square of the concentration, producing an upward curving line in the rate-concentration plot. For this type of reaction, the rate of reaction decreases rapidly (faster than linearly) as the concentration of the reactant decreases.

=CATALYSTS SPEED IT UP= A catalyst is like adding a bit of magic to a reaction. Reactions need a certain amount of energy to happen. If they don't have it, oh well, the reaction probably can't happen. A catalyst lowers the amount of energy needed so that a reaction can happen easier. A catalyst is about energy; it doesn't have to be another molecule. If you fill a room with hydrogen gas and oxygen gas, very little will happen. If you light a match in that room (or just a spark), all of the hydrogen and oxygen will combine to create water molecules. It is an explosive reaction.

The energy needed to make a reaction happen is called the **activation energy**. As everything moves around, energy is needed. The energy a reaction needs is usually in the form of heat. When a catalyst is added, something special happens. Maybe a molecule shifts it's structure. Maybe that catalyst makes two molecules combine and they release a ton of energy. That extra energy might help another reaction to occur. In our earlier example, the spark added the **activation energy**.

Catalysts are also used in the human body, not to cause explosions but to make very difficult reactions happen. They help very large molecules combine. There is another interesting fact about catalysts. Catalysts lower the activation energy required for a reaction to occur. With the activation energy lower, the products can also combine more easily. Therefore, the forward and reverse reactions are both accelerated. It helps both reactions.

=INHIBITORS SLOW IT DOWN= There is also something called an inhibitor that works exactly the opposite of catalysts. Inhibitors slow the rate of reaction. Sometimes they even stop the reaction completely. You might be asking, "Why would anyone need those?" You could use an inhibitor to make the reaction slower and more controllable. Without them, some reactions could keep going and going and going. If they did, all of the molecules would be used up. That would be bad, especially in your body.

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