It’s Friday night, and you’re on the first call of your pediatrics rotation. You’re a third-year medical student now, and you’re anxious to see some action. Hanging out in the call room near the ED of the children’s hospital, you get a page from the resident: Come to the emergency room, now, she says. They just brought in a kid with DKA. DKA, you know, stands for diabetic ketoacidosis and is a fairly common way for undiagnosed Type I diabetes to present. Entering the child’s room, the examination is already under way; he’s young—about 10 years old—conscious but agitated, and the most obvious sign—what you notice immediately—is his rapid, shallow breathing. You see he’s already receiving IV fluids and an insulin drip.
Later that evening, after the boy has stabilized, you and the resident are talking about diabetes and DKA. You remembered from your second-year lessons about endocrine pathophysiology that ketoacidosis can arise as a result of the body’s metabolism of fatty acids when insulin production finally shuts down in Type I diabetes. Because most of the cells of the human body can’t import glucose without the aid of insulin, the glucose accumulates in the plasma of the blood, producing hyperglycemia even as the cells of the body are in a state of glucose starvation. Fatty acids are metabolized into ketone bodies as an alternative energy source. Some of the ketones produced are ketoacids, and as the diabetic crisis continues and worsens, the concentration of these ketoacids increases, resulting in a plasma pH below 7.35 (metabolic acidosis). The combination of the acidosis, progressively severe dehydration due to the osmotic effect of glucose “spilling into” the urine, and other negative effects of the severe insulin depletion result in the host of signs and symptoms of diabetic ketoacidosis. You ask the resident why the boy was hyperventilating, and she takes a piece of paper and writes out the following:
H+ (aq) + HCO3- (aq)
It’s Le Châtelier’s principle, she deadpans, disappointed that you didn’t remember that. The respiratory system is trying to compensate for the metabolic acidosis; the increased breathing rate allows the patient to blow off more CO2, which causes the equilibrium to shift to the right. Hydrogen ions combine with bicarbonate ions to produce carbonic acid, which dissociates into CO2gas to replace the gas that’s being expelled from the lungs. Of course, the desired result is a decrease in the hydrogen ion concentration, which stabilizes the pH and keeps it from getting crazy low. It’s not perfect, but if you catch them soon enough, the pH hasn’t gone so low that they’ve essentially become a scrambled egg. You should know all of this by now.
You recognize the equation. In fact, you even remember studying it for your MCAT. What was that all about? Oh yeah, chemical equilibrium. Wow, chemistry really is essential for medical school!
This chapter focuses on two primary topics: chemical kinetics and chemical equilibrium. As the term suggests, chemical kinetics is the study of reaction rates, the effects of reaction conditions on these rates, and the mechanisms implied by such observations. Chemical equilibrium is a dynamic state of a chemical reaction at which the concentrations of reactants and products stabilize over time in a low-energy configuration. Pay particular attention to the concepts of chemical equilibrium, as we will return to them in our review of solutions, acid-base, and redox chemistry.
Chemical Kinetics
Reactions can be spontaneous or nonspontaneous; the change in Gibbs free energy determines whether or not a reaction will occur, by itself, without outside assistance (see Chapter 6, Thermochemistry). However, even if a reaction is spontaneous, this does not necessarily mean that it will run quickly. In fact, nearly every reaction that our very lives depend upon, while perhaps spontaneous, proceeds so slowly that without the aid of enzymes and other catalysts, we might not ever actually “see” the reaction occur over the course of an average human lifetime. In biology, we discuss the function of enzymes, which selectively enhance the rate of certain reactions (by a factor of 106 to 1014) over other thermodynamically feasible reaction pathways, thereby determining the course of cellular metabolism, the collection of all chemical reactions in a living cell. For now, however, let us review the topics of reaction mechanisms, rates, rate laws, and the factors that affect them.
REACTION MECHANISMS
Very rarely is the balanced reaction equation, with which we work to calculate limiting reactants and yields, an accurate representation of the actual steps involved in the chemical process from reactants to products. Many reactions proceed by more than one step, the series of which is known as the mechanism of a reaction and the sum of which gives the overall reaction (the one that you, typically, are asked to balance). When you know the accepted mechanism of a reaction, this helps you explain the reaction’s rate, position of equilibrium, and thermodynamic characteristics (see Chapter 6). Consider this generic reaction:
Overall reaction: A2 + 2B
Bridge
Mechanisms are proposed pathways for a reaction that must coincide with rate data information from experimental observation. We will be studying mechanisms more in Organic Chemistry.
On its own, this equation seems to imply a mechanism in which two molecules of B collide with one molecule of A2 to form two molecules of AB. Suppose instead, however, that the reaction actually takes place in two steps:
Step 1: A2 + B
Step 2: A2B + B
You’ll note that the two steps, taken together, give the overall (net) reaction. The molecule A2B, which does not appear in the overall reaction, is called an intermediate. Reaction intermediates are often difficult to detect, because they may be consumed almost immediately after they are formed, but a proposed mechanism that includes intermediates can be supported through kinetic experiments. One of the most important points for you to remember is that the slowest step in any proposed mechanism is called the rate-determining step, because it acts like a kinetic “bottleneck,” preventing the overall reaction from proceeding any faster than the slowest step. It holds up the entire process in much the same way that the overall rate of an assembly line production can only be as fast as the slowest step or slowest person (who will probably soon find himself out of a job).
Reaction Rates
Reactions, unfortunately, do not come with handy built-in speedometers. We can’t just look at a dial or gauge and read the reaction rate. It takes a little more effort than that. To determine the rate at which a reaction proceeds, we must take measurements of concentrations of reactants and products and note their change over time.