H2 (g) + Cl2 (g)
Figure 5.2 shows that the reaction is exothermic. The potential energy of the products is less than the potential energy of the reactants; heat is evolved, and the enthalpy change of the reaction is negative.
Figure 5.2
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Kinetics and thermodynamics should be considered separately. Note that the potential energy of the product can be raised or lowered, thereby changing the value of
FACTORS AFFECTING REACTION RATE
If you imagine a group of children playing together in a rough-and-tumble game of tag, we can identify certain conditions that would result in a more vigorous, more fun, more effective game. We can use this analogy to remember the factors that can affect chemical reaction rates.
Reactant Concentrations
A group of children playing tag will probably have more fun the greater the number of children playing: The participation of more children equals more opportunities to chase, trip, and tag each other. The greater the concentrations of the reactants, the greater the number of effective collisions per unit time. Therefore, the reaction rate will increase for all but zero-order reactions. For reactions occurring in the gaseous state, the partial pressures of the gas reactants serve as a measure of concentration (see Chapter 7, The Gas Phase).
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We saw earlier in this chapter that increasing reactant concentrations might increase the reaction rate, but be aware of each reactant’s order before making this assumption.
Temperature
You can imagine that children’s enthusiasm for playing tag outside increases as the temperature warms from the bone-chilling snowstorms of winter to the pleasant sun of June. For nearly all reactions, the reaction rate will increase as the temperature increases. Because the temperature of a substance is a measure of the particles’ average kinetic energy, increasing the temperature increases the average kinetic energy of the molecules. Consequently, the proportion of molecules having energies greater than Ea (thus capable of undergoing reaction) increases with higher temperature. You’ll often hear that raising the temperature of a system by 10°C will result in an approximate doubling of the reaction rate. You have to be careful with this approximation, as it is generally true for biological systems but not so for many other systems. (Don’t forget also: If the temperature gets too high, a catalyst may denature—and then the reaction rate plummets!)
Medium
The rate at which a reaction takes place may also be affected by the medium in which it takes place. Just as children playing tag would prefer to play on a grassy field, but another group of children wanting to get a game of ice hockey going would be looking for an ice rink, some molecules are more likely to react with each other in aqueous environments, while others are more likely to react in a nonaqueous solvent, such as DMSO (dimethylsulfoxide) or ethanol. Furthermore, the physical state of the medium (liquid, solid, or gas) can also have a significant effect. Generally, polar solvents are preferred because their molecular dipole tends to polarize the bonds of the reactants, thereby lengthening and weakening them, which permits the reaction to occur faster.
Catalysts
Catalysts are substances that increase reaction rate without themselves being consumed in the reaction. Catalysts interact with the reactants, either by adsorption or through the formation of intermediates, and stabilize them so as to reduce the energy of activation necessary for the reaction to proceed. While many catalysts, including all enzymes, chemically interact with the reactants, upon formation of the products, they return to their original chemical state. They may increase the frequency of collisions between the reactants; change the relative orientation of the reactants, making a higher percentage of the collisions effective; donate electron density to the reactants; or reduce intramolecular bonding within reactant molecules. In homogeneous catalysis, the catalyst is in the same phase (solid, liquid, gas) as the reactants. In heterogeneous catalysis, the catalyst is in a distinct phase. Figure 5.3 compares the energy profiles of catalyzed and uncatalyzed reactions.
Figure 5.3
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Enzymes are essential to most biological processes.
If you look closely at the energy profiles in the figure, you’ll notice that the only effect of the catalyst is the decrease in the energies of activation, Ea, for both the forward and reverse reactions. The presence of the catalyst has no impact on the potential energies of the reactants or the products or the difference between them. This means that catalysts change only the rate of reactions, and in fact, they change the forward rate and the reverse rate by the same factor. Consequently, they have no impact whatsoever on the equilibrium position or the measure of Keq. Of course, you have to remember that as useful as catalysts are in biological and nonbiological systems, catalysts are not miracle workers: They will not transform a nonspontaneous reaction into a spontaneous one; they only make spontaneous reactions go more quickly toward equilibrium.
Equilibrium
We’ve been dancing around this term for the past couple of pages now. We warned you not to confuse the chemical equilibrium expression for the rate expression. We stressed that catalysts make reactions go faster toward their equilibrium position but that they can’t actually change the equilibrium position or alter the value of Keq. Well, now we’re really going to get into it—and you’d better start paying attention, because the principles and concepts that are the focus of the rest of this chapter will direct our discussion in the upcoming chapters about some of the most important general chemistry topics for the MCAT: solutions, acids and bases, and redox reactions.
DYNAMIC EQUILIBRIUM OF REVERSIBLE CHEMICAL REACTIONS
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Equilibrium, like biological homeostasis, is a dynamic process that seeks to find balance in all systems. We can use this concept to our advantage on the MCAT in all four of the basic sciences. Equilibria are dynamic, meaning that they do undergo change but their net change will be zero.
So far, in our discussion of reaction rates, we have been assuming that the reactions were irreversible; that is, the reaction proceeds in one direction only, the reaction goes to completion, and the amount of product formed is the maximum as determined by the amount of limiting reactant present. Reversible reactions are those in which the reaction can proceed in one of two ways: forward and reverse. (From the perspective of the direction in which the overall reaction is written, the forward reaction is the one that goes from “reactants” on the left to “products” on the right.) Reversible reactions usually do not proceed to completion because (by definition) the products can react together to re-form the reactants. When the reaction system is closed and no products or reactants are removed or added, the system will eventually “settle” into a state in which the rate of the forward reaction equals the rate of the reverse reaction and the concentrations of the products and reactants are constant. In this dynamic equilibrium state, the forward and reverse reactions are occurring—they haven’t stopped, as would be the case in a static equilibrium—but they are going at the same rate; thus, there is no net change in the concentrations of the products or reactants. Consider the following generic reversible reaction: