Getting Started
Chemical kinetics is the branch of chemistry that explores the speed, or rate, at which chemical reactions occur. At the macroscopic level, we can observe some reactions happening in an instant, like an explosion, while others, like the rusting of iron, take years. The core problem of kinetics is to measure, predict, and control these rates by understanding the molecular-level events that transform reactants into products.
What You Should Be Able to Do
After completing this section, you should be able to:
Define reaction rate in terms of the change in concentration of reactants or products over time.
Calculate the average rate of a reaction from experimental data.
Relate the rates of consumption of reactants and formation of products using the stoichiometry of a balanced chemical equation.
Explain qualitatively how changes in temperature, concentration, surface area, and the presence of a catalyst affect the rate of a reaction.
Key Concepts & Analysis
The study of reaction rates is fundamentally about dynamics and change. We start with a set of initial conditions and observe how the system evolves over time, paying close attention to the factors that can alter the speed of that evolution.
Baseline Condition: Measuring the Rate of Reaction
A chemical reaction begins with a specific concentration of reactants and, typically, zero products. The reaction rate is defined as the change in concentration of a substance per unit of time. Its most common units are molarity per second (M/s).
For a generic reaction A → B, we can measure the rate in two ways:
By the disappearance of reactant A: As the reaction proceeds, the concentration of A, denoted as
[A], decreases. The change in concentration,Δ[A], will be negative. To make the rate a positive value, we include a negative sign.Rate = -Δ[A] / ΔtBy the appearance of product B: As the reaction proceeds, the concentration of B,
[B], increases. The changeΔ[B]is positive.Rate = Δ[B] / Δt
The rate is not usually constant; it is typically fastest at the beginning of the reaction when reactant concentrations are highest and slows down as reactants are consumed.
The Process: Relating Rates with Stoichiometry
The rates of change for different species in a reaction are not always equal; they are linked by the coefficients in the balanced chemical equation. Consider the synthesis of ammonia:
N₂(g) + 3H₂(g) → 2NH₃(g)
For every one mole of N₂ that reacts, three moles of H₂ react, and two moles of NH₃ are formed. Therefore, their rates of change are related:
Hydrogen (H₂) is consumed three times as fast as nitrogen (N₂).
Ammonia (NH₃) is formed twice as fast as nitrogen (N₂) is consumed.
To define a single, unambiguous rate for the overall reaction, we divide each species' rate by its stoichiometric coefficient:
Rate = -Δ[N₂]/Δt = -(1/3)Δ[H₂]/Δt = +(1/2)Δ[NH₃]/Δt
This ensures that no matter which reactant or product we monitor, we calculate the same value for the overall reaction rate.
The Resulting Change: Factors Influencing Reaction Rate
The baseline rate of a reaction can be deliberately changed by altering the experimental conditions. These changes are best understood through the lens of collision theory, which states that for a reaction to occur, reactant particles must collide with both sufficient energy and the correct orientation.
| Factor | How It Influences the Rate | Resulting Change in Rate |
|---|---|---|
| Reactant Concentration | Increasing the concentration of reactants means there are more particles in a given volume. This leads to more frequent collisions between reactant particles. | Increases. A higher frequency of collisions leads to a higher frequency of successful, reaction-causing collisions. |
| Temperature | Increasing the temperature increases the average kinetic energy of the particles. This has two effects: particles move faster, leading to more frequent collisions, and more importantly, the collisions are more energetic. | Increases. A much larger fraction of the collisions will have enough energy to overcome the activation energy barrier, drastically increasing the rate of successful reactions. |
| Surface Area | For reactions involving solids, breaking a solid into smaller pieces increases the surface area exposed to the other reactants. This allows for more contact points where collisions can occur. | Increases. More available surface area leads to a higher frequency of collisions between the reactants, accelerating the reaction. |
| Catalysts | A catalyst is a substance that increases the reaction rate without being consumed. It does this by providing an alternative reaction pathway, or mechanism, with a lower activation energy (the minimum energy required for a reaction to occur). | Increases. By lowering the energy barrier, a larger fraction of collisions have sufficient energy to be successful, increasing the rate. The catalyst does not change the overall energy released or absorbed (ΔH). |
Key Models & Representations
The following matrix summarizes the primary factors that control reaction rates, connecting them to the principles of collision theory.
| Factor | Mechanism of Action (Collision Theory) | Effect on Rate | Common Example |
|---|---|---|---|
| Concentration | Increases the frequency of particle collisions. | Increases | A wood fire burns faster in a high-oxygen environment than in air. |
| Temperature | Increases both the frequency and, more significantly, the average energy of particle collisions. | Increases | Food cooks faster at a higher oven temperature. |
| Surface Area | Increases the number of exposed particles, leading to more frequent collisions. | Increases | A powdered antacid tablet neutralizes stomach acid faster than a whole tablet. |
| Catalyst | Provides an alternative reaction pathway with a lower activation energy, increasing the fraction of effective collisions. | Increases | Catalytic converters in cars use metals like platinum to speed up the conversion of toxic exhaust gases to safer substances. |
Key Terms, Quantities, & Concepts
Chemical Kinetics: The study of the rates and mechanisms of chemical reactions.
Reaction Rate: The speed at which a chemical reaction proceeds, defined as the change in concentration of a reactant or product per unit time (e.g., M/s).
Collision Theory: A model stating that for a reaction to occur, reactant particles must collide with sufficient energy (activation energy) and in the correct physical orientation.
Activation Energy (Ea): The minimum amount of energy required for colliding reactant particles to transform into products. It represents an energy barrier that must be overcome.
Concentration ([X]): The amount of a substance in a given volume, typically expressed in moles per liter (M), or molarity.
Catalyst: A substance that increases the rate of a chemical reaction by providing an alternative mechanism with a lower activation energy, without being consumed in the overall process.
Stoichiometry: The quantitative relationship between the amounts of reactants and products as defined by the coefficients in a balanced chemical equation.
Skill Snapshots
Causation
Cause: Increasing the temperature of a reaction mixture results in particles having higher kinetic energy, leading to more frequent and more forceful collisions that are more likely to overcome the activation energy.
Cause: Adding a catalyst results in a lower activation energy for the reaction, allowing a greater percentage of molecular collisions to be successful at a given temperature.
Cause: Increasing the concentration of a reactant results in a greater number of particles in the same volume, which increases the probability and frequency of collisions.
Comparison
A reaction at a high temperature proceeds much faster compared to the same reaction at a low temperature because the exponential increase in high-energy collisions is more significant than the linear increase in collision frequency.
The rate of disappearance of a reactant with a stoichiometric coefficient of 2 is twice as fast compared to the rate of appearance of a product with a coefficient of 1.
A catalyzed reaction follows a different, lower-energy pathway whereas an uncatalyzed reaction must proceed over a higher activation energy barrier.
Change and Continuity Over Time (CCOT)
Baseline: At the beginning of a reaction (t=0), the reactant concentrations are at their maximum, and the initial reaction rate is at its fastest.
Change: As the reaction progresses, reactants are consumed and their concentrations decrease, causing the rate of the reaction to slow down.
Change: Concurrently, the concentrations of the products increase from zero.
Continuity: Throughout the reaction, the ratio of the rates of change for all reactants and products remains constant, as determined by the stoichiometry of the balanced equation.
Common Misconceptions & Clarifications
Misconception: The rate of reaction is the same as the rate of disappearance for any given reactant.
Clarification: The rates are only the same if the stoichiometric coefficient of the reactant is 1. For the reaction
2A → B, reactant A disappears twice as fast as product B appears. The overall reaction rate is defined by taking stoichiometry into account:Rate = - (1/2)Δ[A]/Δt = +Δ[B]/Δt.Misconception: Catalysts are consumed during a reaction.
Clarification: Catalysts actively participate in the reaction by forming temporary intermediates, but they are regenerated in a later step. Therefore, there is no net consumption of the catalyst over the course of the reaction.
Misconception: A fast reaction is always a "good" or "complete" reaction.
Clarification: Reaction rate (kinetics) is entirely separate from reaction yield or extent (equilibrium and thermodynamics). A reaction can be extremely fast but only convert a tiny fraction of reactants to products before reaching equilibrium.
Misconception: Increasing temperature speeds up reactions simply by making molecules move faster and collide more often.
Clarification: While increased collision frequency plays a role, it is a minor one. The primary reason is the significant increase in the energy of the collisions. The fraction of molecules possessing energy equal to or greater than the activation energy increases exponentially with temperature, leading to a dramatic rise in the rate.
One-Paragraph Summary
Chemical kinetics quantifies the speed of reactions by measuring the rate of change in reactant or product concentrations over time. This rate is not static; it is fastest at the start and slows as reactants are depleted. The stoichiometry of a balanced equation provides the precise mathematical relationship between the rates of consumption and formation of all chemical species involved. Critically, the reaction rate can be controlled by manipulating experimental parameters. Increasing reactant concentration or surface area enhances the frequency of collisions, while increasing temperature boosts both the frequency and, more importantly, the energy of these collisions. Catalysts provide an entirely different, lower-energy pathway, accelerating the reaction without being consumed.