Java Security Settings:
      This web page employs Java, which requires specific security settings for correct operation.
      If the applets on this page do not run correctly, consult the Virtual Chemistry Experiments FAQ
      or the Physlet Physics web site for establishing the correct security settings.

Chemical Kinetics

Differential Rate Laws


In many reactions, the rate of reaction changes as the reaction progresses. Initially the rate of reaction is relatively large, while at very long times the rate of reaction decreases to zero (at which point the reaction is complete). In order to characterize the kinetic behavior of a reaction, it is desirable to determine how the rate of reaction varies as the reaction progresses.

A rate law is a mathematical equation that describes the progress of the reaction. In general, rate laws must be determined experimentally. Unless a reaction is an elementary reaction, it is not possible to predict the rate law from the overall chemical equation. There are two forms of a rate law for chemical kinetics: the differential rate law and the integrated rate law.

The differential rate law relates the rate of reaction to the concentrations of the various species in the system.

Differential rate laws can take on many different forms, especially for complicated chemical reactions. However, most chemical reactions obey one of three differential rate laws. Each rate law contains a constant, k, called the rate constant. The units for the rate constant depend upon the rate law, because the rate always has units of mole L-1 sec-1 and the concentration always has units of mole L-1.

Zero-Order Reaction

For a zero-order reaction, the rate of reaction is a constant. When the limiting reactant is completely consumed, the reaction abrupts stops.

Differential Rate Law:       r = k

The rate constant, k, has units of mole L-1 sec-1.

First-Order Reaction

For a first-order reaction, the rate of reaction is directly proportional to the concentration of one of the reactants.

Differential Rate Law:       r = k [A]

The rate constant, k, has units of sec-1.

Second-Order Reaction

For a second-order reaction, the rate of reaction is directly proportional to the square of the concentration of one of the reactants.

Differential Rate Law:       r = k [A]2

The rate constant, k, has units of L mole-1 sec-1.

These three behaviors are illustrated in the following plots. The graph at the left shows concentration-time plots for zero-order (red line), first-order (green line), and second-order (blue line) reactions. The corresponding rate-concentration plots are shown at the right.

In examining the plots, bear in mind that as the reaction progresses, the concentration of reactant decreases. This corresponds to moving from right to left on the plot of reaction rate vs concentration. In this example, the reactant has a stoichiometric coefficient of one, so the reaction rate (plotted in the graph at the right) corresponds with the negative value of the slope of the concentration-time curve (plotted in the graph at the left). Carefully examine the graphs and take note of the following points:



Consider the following reaction between formic acid (HCOOH) and bromine:

HCOOH (aq) + Br2 (aq)   →   2 H+ (aq) + 2 Br- (aq) + CO2 (aq)

In this chemical system, the only species that absorbs visible light is bromine; thus spectrophotometry can be employed to determine how the concentration of bromine varies with time during the reaction.

One syringe contains a 2.00 mM solution of bromine and the other syringe contains a 0.200 M solution of formic acid. The two liquids are mixed in equal portions. As the reaction is occurring, the concentration of bromine is plotted versus time in the left graph and the rate of reaction is plotted versus the bromine concentration in the right graph.

Examine the two graphs and determine the differential rate law for this reaction. Calculate the rate constant using values read from the graph(s).

Rate of Reaction                     Integrated Rate Laws

Chemical Kinetics Home Page
Virtual Chemistry Home Page

DifferentialRateLaws.html version 2.1
© 2000-2014 David N. Blauch