The Pharmaceutics and Compounding Laboratory
Chemical Kinetics


Hydrolysis of the drug entity can be a major factor in the instability of solutions. Aspirin, for example, undergoes hydrolysis with the resultant degradation products being salicylic acid and acetic acid. The rate of this reaction is said to be second order, since it is dependent not only upon the aspirin concentration, but upon solution pH (i.e. the hydronium ion concentration at solution pH values less than approximately 2.5 or the concentration of hydroxyl ion at solution pH values greater than approximately 7.0). At pH = 7.5, the rate expression for the hydrolysis of aspirin may be written:


[A] = the concentration of aspirin - [OH-] = the hydroxyl ion concentration - K = the second order rate constant - t = time

If the solution is buffered so that the hydroxyl ion concentration remains essentially constant, the rate expression may be rewritten as:


C = the unchanging hydroxyl ion concentration

Since two constants can always be combined into one constant, the above expression is equal to:


Kapp = KC

From the above equation, it can be seen that the degradation of aspirin in a solution buffered at pH = 7.5 will follow first order kinetics; that is, the reaction will appear to be a first order reaction, dependent only on the concentration of one reactant; i.e. aspirin.

The integrated form of a first order rate expression is:


At = the amount of drug remaining at time = t - Ao = the amount of drug initially present - Kapp = the apparent first order rate constant - t = time of sampling

This equation is of the form:

y = mx + b


m = the slope of the line - b = the y intercept

For the hydrolysis of aspirin in buffered solution (pH = 7.5), a semi-log plot of the aspirin concentration remaining versus time should yield a straight line with a negative slope equal to -Kapp.

The experimentally determined first order rate constant (Kapp) can be related to the true second order rate constant by the expression:

Kapp = K[OH-]

The pseudo first order degradation of aspirin in a solution buffered at pH = 7.5 can be followed by measuring the increasing concentration of salicylic acid spectrophotometrically.

One mole of salicylic acid is produced when one mole of aspirin degrades; so, using the ratio of the molecular weights of aspirin to salicylic acid, we can determine the weight of aspirin degraded for each mg of salicylic acid produced.

Thus, each milligram of salicylic acid present represents the degradation of 1.304 milligrams of aspirin. Since the amount of aspirin initially present is known and since the amount of aspirin which has degraded can be determined, the amount of aspirin remaining can be calculated.