Multistep energy profiles provide a visual way to connect thermodynamic “ups and downs” to kinetics. By comparing the relative heights of the barriers between intermediates and transition states, you can often identify which step most strongly limits the overall reaction rate.

Interpreting “barriers” on a multistep energy profile

A multistep reaction energy profile shows several peaks (transition states) and valleys (reactants/intermediates/products).

A two-step reaction coordinate diagram showing reactants, an intermediate, and two transition states (two peaks). The key kinetic comparison is the vertical gap from each valley to the next peak (each step’s $E_a$), not the absolute height of a peak relative to the original reactants. The step with the larger valley-to-peak barrier is typically the slow (rate-determining) step. Source

Each elementary step has its own activation barrier measured from its starting valley to its peak.

When the syllabus says “compare barrier heights,” it means comparing the vertical energy gaps for each step, not simply comparing absolute peak heights measured from the original reactants.

How to compare barriers correctly

• Identify each local minimum (reactants or intermediates).

• For each step, locate the next transition state peak immediately “ahead” along the reaction coordinate.

• Compute (conceptually) each step’s barrier as: peak height minus the preceding valley height.

• The largest barrier (largest $E_a$) is the step that is usually the slowest.

Why the largest barrier usually corresponds to the slow step

A higher activation barrier means a smaller fraction of collisions/encounters have enough energy to reach the transition state at a given temperature, so that step proceeds less frequently per unit time.

This connection is often justified by the Arrhenius relationship between rate constant and activation energy.

An Arrhenius plot of $\ln(k)$ versus $1/T$ illustrating the linear form of the Arrhenius equation. The line’s slope is proportional to $-E_a/R$, so a larger activation energy corresponds to a steeper negative slope and a stronger temperature dependence of the rate constant. The intercept relates to the pre-exponential factor $A$. Source

Even without calculation, the key idea is: larger $E_a$ ⇒ exponentially smaller $k$ ⇒ slower step, so the largest barrier commonly controls the overall timescale.

Connecting the “largest barrier” to what you observe experimentally

In many multistep profiles:

• The highest barrier corresponds to the smallest step rate constant.

• Intermediates before that barrier may accumulate relative to others, because they form faster than they are consumed.

• The overall reaction appears to “wait” for that difficult barrier-crossing event.

Forward vs reverse direction matters

Barrier comparisons must be made in the direction of interest:

• The forward slow step is linked to the largest forward barrier (from each valley to the next peak).

• If you considered the reverse reaction, you would compare barriers from products/intermediates toward reactants; the largest reverse barrier might be different.

Common pitfalls when using energy profiles to identify the slow step

• Comparing absolute peak heights instead of valley-to-peak gaps for each step.

• Assuming the tallest peak from the baseline reactants is always rate-limiting, even when there are high-energy intermediates later.

• Ignoring that a step can be “large barrier” yet not control the observed rate if it is not on the dominant pathway or if conditions change (temperature can change the relative importance of barriers).

• Confusing overall energy change (reactants vs products) with kinetics; being more exothermic or endothermic does not, by itself, determine which step is slow.