Reaction Rate
The reaction rate is the change in concentration of a reactant or product per unit time in a chemical reaction, expressed in units of mol/(L·s) or mol·L⁻¹·s⁻¹. It quantifies how quickly reactants are consumed and products are formed, and is influenced by factors including concentration, temperature, surface area, catalysts, and the nature of the reactants. Understanding reaction rates is fundamental to chemical engineering (designing reactors), pharmacology (drug metabolism), and environmental chemistry (pollutant breakdown).
Key Formula
rate = −(1/a) × Δ[A]/Δt = +(1/b) × Δ[B]/Δt
LaTeX: r = -\frac{1}{a}\frac{\Delta[A]}{\Delta t} = +\frac{1}{b}\frac{\Delta[B]}{\Delta t}
| Symbol | Meaning | Unit |
|---|---|---|
| r | Reaction rate | mol/(L·s) |
| [A] | Molar concentration of reactant A | mol/L |
| [B] | Molar concentration of product B | mol/L |
| a, b | Stoichiometric coefficients | dimensionless |
| t | Time | s |
Worked Example
Problem
In the reaction 2NO(g) + O₂(g) → 2NO₂(g), the concentration of NO decreases from 0.100 mol/L to 0.068 mol/L in 10.0 seconds. Calculate the average rate of reaction and the rate of formation of NO₂.
Solution
Step 1: Calculate the rate of disappearance of NO: Δ[NO] = 0.068 − 0.100 = −0.032 mol/L Δt = 10.0 s Step 2: Rate with respect to NO: rate = −(1/2) × (Δ[NO]/Δt) = −(1/2) × (−0.032/10.0) = −(1/2) × (−3.2 × 10⁻³) = 1.6 × 10⁻³ mol/(L·s) Step 3: Rate of formation of NO₂: Since stoichiometry of NO₂ matches NO (both coefficient 2): Δ[NO₂]/Δt = −Δ[NO]/Δt = +3.2 × 10⁻³ mol/(L·s)
Answer
Average reaction rate = 1.6 × 10⁻³ mol/(L·s); Rate of NO₂ formation = 3.2 × 10⁻³ mol/(L·s)
Factors Affecting Reaction Rate
| Factor | Effect on Rate | Explanation | Example |
|---|---|---|---|
| Concentration | Increases with concentration | More collisions per unit time | Higher [HCl] dissolves Zn faster |
| Temperature | Rate doubles per ~10°C rise | More molecules exceed Eₐ | Food spoils faster when warm |
| Surface area | Increases with finer particles | More exposed reaction sites | Powdered Zn reacts faster than lumps |
| Catalyst | Increases (lowers Eₐ) | Provides alternate reaction pathway | MnO₂ catalyzes H₂O₂ decomposition |
| Nature of reactants | Varies by bond type | Ionic bonds react faster than covalent | NaCl dissolves fast; diamond does not |
Interactive Tools
Wikimedia Commons, CC BY-SA
Related Terms
Rate Law
The rate law (or rate equation) is a mathematical expression that relates the reaction rate to the concentrations of reactants raised to experimentally determined powers called reaction orders. It takes the general form rate = k[A]^m[B]^n, where k is the rate constant and m, n are the orders with respect to each reactant. Critically, the rate law cannot be deduced from the stoichiometric equation but must be determined experimentally, and it reflects the mechanism of the rate-determining step.
Activation Energy
Activation energy (Eₐ) is the minimum amount of energy that reacting molecules must possess for a collision to result in a chemical reaction — effectively the energy barrier that must be overcome to convert reactants into products. It determines how fast a reaction proceeds: reactions with low activation energies are generally fast (explosions, acid-base), while those with high activation energies are slow (rusting, digestion). The concept was introduced by Svante Arrhenius in 1889 and is central to the Arrhenius equation and transition state theory.
Chemical Catalysis
Chemical catalysis is the process by which a catalyst — a substance that participates in a reaction and increases its rate without being consumed or permanently altered — provides an alternative reaction pathway with a lower activation energy. Catalysts can be homogeneous (same phase as reactants, e.g., H⁺ in acid hydrolysis), heterogeneous (different phase, e.g., Pt in catalytic converters), or biological (enzymes). Catalysis is fundamental to industrial chemistry: approximately 85-90% of all industrial chemical processes rely on catalysts, including the Haber-Bosch ammonia synthesis (Fe catalyst) and petroleum cracking (zeolites).
From Latin "reactio" (a backward action, response) and Old English "raet" (a reckoning). In chemistry, "reaction" was used from the 17th century to describe chemical transformations, and "rate" entered the scientific vocabulary in the 19th century when quantitative kinetics developed.