Average Reaction Rate Calculator

At a glance:The reaction rate is a physical quantity measuring how fast a chemical reaction proceeds — defined as the amount of reactant consumed or product formed per unit time. The most common and the one this calculator uses is the 'average reaction rate', describing the average concentration change over a time interval: average rate = |C₂ − C₁| / Δt = |ΔC| / Δt, where C₁, C₂ are the initial and final concentrations (usually mol/L, i.e. M) and Δt the elapsed time (usually seconds). We take the absolute value of the concentration change because a reactant's concentration 'decreases' (ΔC negative) while a product's 'increases' (ΔC positive); to keep the rate always positive for comparison we take the absolute value and state which species and whether it is consumed or formed. Using this tool's default: a reactant's concentration falls from 0.5 mol/L to 0.2 mol/L over 10 s, giving average rate = |0.2 − 0.5| / 10 = 0.3 / 10 = 0.03 mol/(L·s). Distinguish 'average rate' from 'instantaneous rate': average rate is the total change over a period divided by the time, reflecting the overall average of that interval; instantaneous rate is the rate at a specific moment, equal to the slope of the concentration–time curve at that point (dC/dt in calculus). Since most rates slow as the reaction proceeds (reactants deplete), the average rate usually lies between the instantaneous rates at the interval's ends; the shorter the interval, the closer the average rate is to the instantaneous. Moreover, if a stoichiometric coefficient is not 1, divide by it: for aA + bB → cC + dD, the unified rate = −(1/a)·d[A]/dt = −(1/b)·d[B]/dt = (1/c)·d[C]/dt = (1/d)·d[D]/dt. This calculator gives the rate 'from a chosen species' concentration change; divide by that species' coefficient for the unified rate. Five main factors affect rate: (1) reactant concentration — higher concentration, more frequent collisions, usually faster (basis of rate law rate = k[A]^m[B]^n); (2) temperature — higher temperature increases molecular kinetic energy and effective collisions, greatly raising rate (Arrhenius equation; the rule of thumb 'rate roughly doubles per 10°C rise'); (3) catalyst — lowers activation energy, providing a new path, accelerating without being consumed; (4) surface area — finer solid particles, larger area, faster; (5) pressure (gas reactions) — higher pressure is like higher concentration, faster. Reaction rate belongs to 'chemical kinetics', a different level from 'chemical thermodynamics' (e.g. Gibbs free energy) which judges spontaneity: thermodynamics tells whether a reaction 'can happen spontaneously', kinetics tells 'how fast'. Notes: first, C₁, C₂ units must match (both mol/L). Second, Δt must be > 0 (this tool returns 0 otherwise). Third, the result is an average rate, not an instantaneous one. Fourth, divide by the stoichiometric coefficient for the unified rate.