On a reaction coordinate diagram (also called a potential energy diagram), the catalyzed reaction is the curve with the lower peak. Both curves start at the same energy level (reactants) and end at the same energy level (products), but the catalyzed curve rises to a noticeably shorter maximum in between. That shorter peak represents a lower activation energy, which is the defining visual signature of a catalyzed reaction.
How to Spot the Catalyzed Curve
A typical textbook diagram shows two curves overlaid on the same axes. The x-axis is the “reaction coordinate” (essentially the progress of the reaction from start to finish), and the y-axis is potential energy. Both curves begin at the same point on the left (reactant energy) and end at the same point on the right (product energy). The only difference is what happens in the middle.
The uncatalyzed curve has a tall, smooth hump. The catalyzed curve has a shorter hump, sometimes split into two smaller peaks with a shallow valley between them. That valley represents a short-lived intermediate species that forms during the alternate pathway the catalyst provides. If you see two smaller bumps instead of one large one, that’s the catalyzed pathway showing a two-step mechanism with two transition states.
The key measurement is activation energy: the vertical distance from the reactant energy level up to the highest point on the curve. Whichever curve has the smaller activation energy is the catalyzed reaction. In some textbook examples, the enzyme-catalyzed version of a biochemical reaction requires roughly one-third the activation energy of the uncatalyzed version.
Why the Peak Is Lower
A catalyst works by offering the reaction a completely different route from reactants to products. Instead of forcing molecules over one tall energy barrier, it breaks the process into smaller steps, each with a lower barrier. The catalyst participates in these intermediate steps but is neither created nor destroyed, so it doesn’t appear in the balanced chemical equation.
Think of it like a mountain pass. The uncatalyzed reaction forces molecules to climb straight over the summit. The catalyzed reaction routes them through a lower pass that may wind through a valley on the way. The starting elevation and ending elevation are the same either way, but the highest point along the catalyzed path is significantly lower.
What Stays the Same on Both Curves
A catalyst does not change the energy of the reactants, the energy of the products, or the overall energy change of the reaction. The enthalpy change (the difference between product energy and reactant energy) is identical for both curves. This means a catalyst cannot make an energetically unfavorable reaction suddenly favorable. It only makes the reaction happen faster by reducing the energy barrier molecules need to overcome.
On the diagram, this is easy to confirm. The left side of both curves sits at the same height, and the right side of both curves sits at the same height. If those values differed, the diagram would be drawn incorrectly.
Why a Lower Peak Means a Faster Reaction
The connection between a lower peak on the diagram and a faster reaction comes down to probability. At any given temperature, molecules have a range of kinetic energies. Only the fraction of molecules with enough energy to clear the activation energy barrier can actually react. When a catalyst lowers that barrier, a much larger fraction of molecules qualify, so collisions lead to products far more often.
This relationship is captured mathematically by the Arrhenius equation, where the reaction rate constant depends on the term e raised to the power of negative activation energy divided by temperature. Because activation energy sits in a negative exponent, even a modest decrease produces a large increase in the rate constant. A catalyst that cuts activation energy by a third, for example, doesn’t just speed the reaction up by a third. It can increase the rate by orders of magnitude, depending on the temperature.
Common Diagram Variations
Not every diagram labels the curves explicitly, which is why knowing the visual cues matters. Here are the most common formats you’ll encounter:
- Two separate smooth curves: The taller one is uncatalyzed, the shorter one is catalyzed. Both share the same start and end points.
- One smooth curve and one with two humps: The two-humped (lower) curve is catalyzed, showing a two-step mechanism with an intermediate. The smooth, taller curve is uncatalyzed.
- Labeled as (a) and (b): Some diagrams place the uncatalyzed reaction in panel (a) and the catalyzed reaction in panel (b). Compare the height of the transition state (the peak) relative to the reactant energy. The panel with the lower transition state is the catalyzed reaction.
Regardless of format, the rule never changes: the curve with the lower activation energy, meaning the smaller vertical gap between reactants and the highest peak, is always the catalyzed reaction.