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How does increasing reactant concentration affect a chemical reaction rate?

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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.

Increasing the concentration of reactants generally leads to an increase in the rate of a chemical reaction.1

According to collision theory, a reaction occurs when reactant particles collide with sufficient activation energy and proper orientation. By increasing the concentration of reactants, you increase the number of particles present in a given volume, which raises the frequency of these collisions. Because the reaction rate is determined by the frequency of effective collisions, a higher concentration of reactants typically results in a faster reaction.2

In chemical kinetics, the relationship between the rate of a reaction and the concentrations of the reacting substances is expressed through the rate constant. While the historical law of mass action suggested that the rate is proportional to the product of the masses of the reactants, modern chemistry recognizes that this relationship is specific to the mechanism of the reaction.13