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Why do chemical reaction rates usually slow down as reactants are used up?

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A graph showing a potential energy curve for a chemical reaction, with the y-axis labeled potential energy and the x-axis labeled progress of reaction. The curve rises from the initial state (reactants) to a peak labeled transition state (activated complex), indicating activation energy, then drops to the final state (products). — chemical kinetics In the Arrhenius equation, the activation energy (E) represents the minimum amount of energy required to transform reactants into products in a chemical reaction. On a potential energy curve, the value of the activation energy is equivalent to the difference in potential energy between particles in an intermediate configuration (known as the activated complex, or transition state) and particles of reactants in their initial state. The activation energy thus can be visualized as a barrier that must be overcome by reactants before products can be formed.

In chemical kinetics, the rate of a reaction is generally dependent on the concentration of the reacting substances. As a reaction progresses, reactants are consumed, leading to a decrease in their concentration over time.123

According to the collision theory, a chemical reaction occurs when molecules collide with sufficient energy and in the correct orientation. Because the reaction rate is proportional to the concentration of the reactants, a reduction in the number of reactant molecules per unit volume results in fewer successful collisions per unit of time. Consequently, as the reactants are depleted, the frequency of these effective collisions decreases, causing the overall reaction rate to slow down.412

While this is the typical behavior for most reactions, it is important to note that for zeroth-order reactions, the rate remains constant regardless of the concentration of the reacting substance.2