12.3 Rate laws  (Page 4/10)

Rate laws may exhibit fractional orders for some reactants, and negative reaction orders are sometimes observed when an increase in the concentration of one reactant causes a decrease in reaction rate. A few examples illustrating these points are provided:

It is important to note that rate laws are determined by experiment only and are not reliably predicted by reaction stoichiometry.

Reaction orders also play a role in determining the units for the rate constant k . In [link] , a second-order reaction, we found the units for k to be $\text{L}{\text{mol}}^{\mathrm{-4}}{\text{s}}^{\mathrm{-1}},$ whereas in [link] , a third order reaction, we found the units for k to be mol ⁻² L ² /s. More generally speaking, the units for the rate constant for a reaction of order $\left(m+n\right)$ are ${\text{mol}}^{1\text{−}\left(m\text{+}n\right)}{\text{L}}^{\left(m\text{+}n\right)\mathrm{-1}}{\text{s}}^{\mathrm{-1}}.$ [link] summarizes the rate constant units for common reaction orders.

Note that the units in the table can also be expressed in terms of molarity ( M ) instead of mol/L. Also, units of time other than the second (such as minutes, hours, days) may be used, depending on the situation.

Key concepts and summary

Rate laws provide a mathematical description of how changes in the amount of a substance affect the rate of a chemical reaction. Rate laws are determined experimentally and cannot be predicted by reaction stoichiometry. The order of reaction describes how much a change in the amount of each substance affects the overall rate, and the overall order of a reaction is the sum of the orders for each substance present in the reaction. Reaction orders are typically first order, second order, or zero order, but fractional and even negative orders are possible.

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