Chemical Kinetics

Isolation Method

A common challenge in chemical kinetics is to determine the rate law for a reaction with multiple reactants. With the concentrations of several species changing simultaneously, real rate laws might not follow simple zero-, first-, and second-order rate laws. The Method of Initial Rates can, in principle, be used to determine the rate law, but in practice the Method of Initial Rates poses serious experimental challenges. The primary challenge is to accurately determine the initial rate.

A better and more commonly used strategy is the Isolation Method. In this method the concentration of one reactant is made much smaller than the concentrations of the other reactants. Under this condition, all reactant concentrations except one are essentially constant, and the simple zero-, first-, and second-order kinetic plots often be used to interpret the concentration-time data.

To see how this method works, consider the following substitution reaction:

CH3Br (aq)   +   OH- (aq)   →   CH3OH (aq)   +   Br- (aq)

The rate law for this substitution reaction is first-order in bromomethane and first-order in hydroxide ion; second-order overall.

r = k [CH3Br] [OH-]

Suppose the reaction were performed starting with 0.100 mole/L OH- and 0.00100 mole/L CH3Br. When the reaction is finished, the solution will contain 0.099 mole/L OH- and essentially no CH3Br. Notice that the concentration of hydroxide ion is approximately unchanged. The total change in [OH-] is 1%, which is often comparable to or smaller than the experimental error. The value of [OH-] can therefore be treated as a constant, and the rate law becomes

r = kobs [CH3Br]

The above rate law is first-order. The observed rate constant, kobs, is a pseudo-first-order rate constant. The term pseudo rate constant means kobs is not the real rate constant. The reaction only appears to be first order with a rate constant kobs. In reality, the reaction is second order overall and the rate constant is k. The pseudo-first-order rate constant is related to the true rate constant by

kobs = k [OH-]

Performing the reaction with 0.100 mole/L OH- and 0.00100 mole/L CH3Br allows one to determine the order of the reaction with respect to CH3Br by creating the characteristic graphs for various rate laws and looking for the graph that is linear. To determine the order of the reaction with respect to OH-, one would perform a series of experiments in which [OH-] varies but is always much larger than [CH3Br]. For each experiment, kobs would be determined, and one would then prepare a plot of ln kobs vs ln [OH-]. Such a plot should be linear with a slope of b.

kobs = k [OH-]b

ln kobs = ln k + b ln [OH-]

Once the value of b is known, one calculate a value for k for each experiment. These values for k would be averaged to obtain the best estimate for the rate constant.

How large a concentration excess is required to effectively "isolate" the effects of a single reactant? As a general rule, a minimum of a 20-fold stoichiometric excess is necessary. A 50-fold or 100-fold stoichiometric excess is preferrable.




An important reaction that occurs in the atmosphere is

2 NO (g)   +   O2 (g)   →   2 NO2 (g)

Nitrogen oxides (NOx) are produced by automobile engines, and the fate of these oxides in the environment is of considerable concern. Modeling reactions involving nitric oxide (NO) and nitrogen dioxide (NO2) are important in understanding the behavior of air pollution and in developing effective regulations. Nitric oxide is a reactive, colorless gas; nitrogen dioxide is an orange-brown gas that is responsible for the characteristic color of Los Angeles smog.

In this experiment, the stopped-flow apparatus has been adapted to handle gases. Both syringes are charged with gas mixtures at 1 atm pressure. The syringe on the left contains 99% N2 and 1% NO, which at 25o C and 1 atm corresponds with 4.09 x 10-4 mole/L NO. The syringe on the right contains pure oxygen, which corresponds to 0.0409 mole/L O2. Both syringes are loaded with 10 mL of gas.

The reactants are colorless and the NO2 product has an orange-brown color. Thus the NO2 concentration is readily measured experimentally using visible spectrophotometry. The graph at the right of the stopped-flow apparatus shows how the concentration of NO2 varies with time during the experiment. (This is the characteristic zero-order kinetic plot.) By clicking the appropriate button, characteristic first-order and second-order kinetic plots can also be prepared. In the first- and second-order plots, it is necessary to plot [NO2]f - [NO2] rather than [NO2], because NO2 is a product rather than a limiting reactant. [NO2]f is the final nitrogen dioxide concentration, which is known from the initial concentrations and the reaction stoichiometry.

Perform this reaction several times, using various proportions of NO and O2. Note that for each volume proportion available for the experiment, the concentration of O2 is always present in large excess. The concentration of O2 never changes by more than 1% during any of the reaction conditions. For each experiment, use the appropriate kinetic plot to determine the order of the reaction with respect to NO and the pseudo-rate constant kobs for the chosen oxygen concentration.

For each experiment, enter the oxygen concentration and kobs in the table at the bottom of the page. The points in the table are used to construct plot of ln kobs vs ln [O2]. Use the slope of this plot to determine the order of the reaction with respect to oxygen. Then go back to the set of values for [O2] and kobs and calculate a value for k for each trial. Finally obtain an average value for k.

Here is a list of things to remember when analyzing the experimental data.


Volumes Delivered     mL NO : mL O2

Kinetic Plots

[NO2] vs Time



L mole-1 sec-1

ln [O2]
ln kobs


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© 2000, 2014, 2023 David N. Blauch