What is a zero-order reaction?

A zero-order reaction is one whose rate is independent of reactant concentration—doubling [A] doesn’t change the rate. This behavior often appears when a surface or catalyst is saturated (e.g., a metal surface has all active sites occupied) or in enzyme systems at very high substrate concentration. The integrated rate law is linear: [A]ₜ = [A]₀ − k t, so a plot of concentration versus time is a straight line with slope −k. The half-life depends on starting concentration: t₁/₂ = [A]₀/(2k). Units of k are concentration·time⁻¹ (e.g., mol L⁻¹ s⁻¹), reflecting the constant rate. 

To diagnose zero-order behavior in class or the lab, plot [A] vs t for your data; a straight line indicates zero order. If not, try ln[A] vs t (first order) or 1/[A] vs t (second order). Real-world example: the decomposition of ammonia on a tungsten surface can show zero-order kinetics when the surface is saturated.

Actionable tip: when you encounter zero-order kinetics in a process, reducing initial concentration won’t slow the initial rate—only reducing k (e.g., lowering temperature or poisoning active sites) will.