Chemical Equilibrium

A catalyst is something that will change the rate of the reaction by decreasing the activation energy. It will not change the energy of the reactants or the energy of the end products but will only change the activation energy. Figure 31 describes what this looks like:

Figure 31.

What a catalyst does is modify the transition state energy so that the overall activation energy is lowered. This is what’s seen in enzymes that make it more likely to have the transition state occur. Catalysts will not be consumed as part of the reaction so that they will continue to drive the reaction forward. Because a catalyst does not get consumed, small amounts of any given catalyst are necessary in order to drive the chemical reaction.

CHEMICAL EQUILIBRIUM

So far, we’ve talked about reactions as though they go from one point to another point but do not go back and do not stop until the limiting reactant has been used up. The concept of chemical equilibrium involves the idea that, at the time of the end of the reaction, there will be equality in the rate of the forward reaction and in the rate of the

reverse direction. So, when describing a reaction, the convention of using a forward arrow to represent the reaction is instead a forward and reverse arrow, which takes into account that the equilibrium state has been achieved. It can only be achieved in a closed system.

Catalysts will decrease the activation energy but will have no effect on the equilibrium point. According to Le Chatelier’s principle or the Equilibrium law, a system will have a disturbance in the equilibrium (such as a change in the concentration of reactants or products, temperature, pressure, or volume) and will have a new equilibrium state. There will often be an observable property or measured physical property in the reaction when it has achieved equilibrium. This can be a change in density, color, pressure, or concentration of substances in the system.

As mentioned, the system must be closed in order to have equilibrium. This is because, if one or more of the products (or reactants) escape, there will be a shift of the reaction so that one part of the equation does not participate in the reaction, driving the reaction in one direction. In equilibrium, the process of “reacting” does not cease; rather, there is equivalency of the driving forces of both reactions (forward and reverse).

The chemical reaction in which xA + yB goes to zC, where x, y, and z are the stoichiometric constants in the reaction and A, B, and C, refer to the concentration of the different reactants and products, the equilibrium constant can be referred to using Figure 32:

Figure 32.

Note that the coefficients of products and reactants are powers of the concentration values. While the concentrations of the products and reactants may change, the ratio of these concentrations will be the same so that the value of K will remain constant. This is why it is called the “equilibrium constant”. If the value of K is greater than 1, the predominance of products will occur; when K is equal to one, neither side is favored;

when K is less than 1, there will be a predominance of reactants. If the K is zero, this implies that no reaction whatsoever is occurring.

There are two types of equilibrium reactions. The first is a homogeneous reaction. This is one in which the states of matter are all the same. It can be liquid, gaseous, or solid but all reactants and products are in the same state. An example of this would be nitrogen gas plus hydrogen gas going to ammonia or NH3. This will ultimately reach a state of equilibrium.

The second is a heterogeneous reaction, in which the states of matter are different from one another. This can happen in an aqueous solution when a precipitate happens or even when a solid is dissolved into aqueous ions. One example, is the decomposition of baking soda or sodium hydrogen bicarbonate (NaHCO3) into sodium carbonate solid, water vapor, and CO2 gas.

So, how do you get to the equilibrium constant? It is not a constant that is a given number that is the same for all reactions. It must be calculated by measuring the concentrations of products and reactants in a system and using the equation in figure 32. The equilibrium constant is called a constant because it determines the rate of a reaction in an ideal reaction that involves a set of reactants and a set of products. When the reaction is gaseous, the concentration is not used; instead, the partial pressures of the different gaseous entities is used in the equilibrium constant equation.

The equilibrium constant works well for gaseous homogeneous reactions and for solutions in water. It is different in heterogeneous reactions. In these, solids and pure liquids are excluded from the equilibrium constant. The activities of solids and liquids are considered to be one for mathematical purposes so they do not affect the overall K value. Solvents also have a mathematical equivalent of one. So, when doing the calculation of K, there is nothing added to the equation by solvents, solids, and liquids.

If there is more than one step in the reaction, each will have its own K value. For mathematical and practical purposes, when there are multiple steps, the K values are not added together but are multiplied together.