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Class 12 · Chemistry

Chemical Kinetics

Chemical kinetics is the science of reaction races. It answers the question: how fast do reactions happen, and what controls their speed?

Feynman Lens

Start with the simplest version: this lesson is about Chemical Kinetics. If you can explain the core idea to a friend using everyday language, examples, and one clear reason why it matters, you have moved from memorising to understanding.

Chemical kinetics is the science of reaction races. It answers the question: how fast do reactions happen, and what controls their speed? Some reactions are explosive (combustion), while others are glacial (diamond formation). The difference isn't thermodynamic—both might be "downhill" energetically—but kinetic: the path and obstacles between reactants and products. This chapter explores how concentration, temperature, and catalysts control reaction rates, and how to predict reaction behavior mathematically.

Reaction Rate: The Speed of Change

Reaction rate is how fast a reaction proceeds, measured as change in concentration per unit time. Imagine a candle burning: the wax mass decreases at a steady rate. The burning is fast at first, then slows as the candle shrinks—average rate versus instantaneous rate.

Average Rate: Total change in concentration divided by total time. Simple but masks what happens during the reaction.

Instantaneous Rate: The speed at a specific moment, found by calculating the slope of the concentration-versus-time curve at that instant. It's always changing as the reaction progresses.

Rate Expression

For a reaction: aA + bB → cC + dD

The rate can be expressed as:

Rate = -Δ[A]/Δt = -Δ[B]/Δt = Δ[C]/Δt = Δ[D]/Δt

The negative signs for reactants indicate concentration is decreasing.

Order of Reaction: The Secret Power Law

Here's the surprise: the rate law isn't determined by the balanced equation. Instead, it comes from experiment. A reaction's order tells you which concentrations matter and how much.

Zero-order reactions: Rate is constant regardless of concentration. Like draining a bathtub—once the drain is open, water flows out at a constant rate. Mathematically: Rate = k (constant).

First-order reactions: Rate depends on concentration to the first power. A radioactive nucleus decays at a rate proportional to how many nuclei remain. Double the nuclei, double the decay rate. Examples: radioactive decay, decomposition of hydrogen peroxide catalyzed by enzymes. Rate = k[A].

Second-order reactions: Rate depends on concentration squared, or on two different concentrations multiplied. This happens when the reaction requires two molecules to collide simultaneously. Rate = k[A]² or k[A][B].

Order vs. Molecularity

Molecularity is the number of molecules that must collide for an elementary reaction step. Order is experimental—determined by watching how rate changes with concentration. They're different concepts. A second-order reaction might occur in a single bimolecular step, or through multiple steps that combine to give second-order behavior.

Activation Energy: The Barrier to Reaction

Think of reactants climbing a mountain to become products. The activation energy (Ea) is the height of that mountain—the minimum energy molecules need to transform. Hydrogen and oxygen sit in a room at room temperature: no explosion. Add heat or a spark (activation), and they explode. The spark provides energy to overcome the barrier.

Rate constant (k) quantifies how temperature affects reaction speed. The Arrhenius equation connects them:

k = A × e^(-Ea/RT)

Where:

As temperature increases, k increases exponentially. A 10°C rise might double or triple the rate—why cakes bake faster in hotter ovens.

Integrated Rate Laws: Predicting Concentration Over Time

Once you know the order, you can predict how concentration changes over time.

Zero-order: [A] = [A]₀ - kt (linear decrease)

First-order: ln[A] = ln[A]₀ - kt (exponential decrease)

Half-life (time for concentration to drop to half) depends on order:

Factors Affecting Reaction Rate

Concentration: More molecules mean more collisions, faster reactions. But the effect depends on order.

Temperature: Higher temperature gives molecules more kinetic energy. They collide harder and more frequently, overcoming activation barriers. A 10°C increase often doubles or triples the rate.

Surface Area: For solid reactants, crushing them into powder increases surface area, allowing more contact with other reactants. Why matches ignite faster when scratched than when held over a flame.

Catalysts: These substances lower activation energy without being consumed. A catalyst provides an alternative pathway over a lower mountain. Homogeneous catalysts (same phase as reactants) mix freely with reactants; heterogeneous catalysts (different phase) work at surfaces. Biological catalysts are enzymes—incredibly selective and efficient.

Nature of Reactants: Ionic reactions in solution are usually fast (charges attract). Reactions between gases are slower than expected (molecules must find and collide). Reactions involving breaking strong bonds are slow.

Reaction Mechanisms: The Molecular Play

The balanced equation shows only start and finish—not the intermediate steps. A reaction mechanism reveals the actual sequence of elementary steps. Consider:

Overall: 2NO + O₂ → 2NO₂

One possible mechanism:

The slow step determines overall rate. The reaction is third-order overall (second in NO, first in O₂) even though the balanced equation suggests second-order in NO. This matches the experimental rate law.

Catalysis: The Reaction Accelerator

Catalysts are the chemist's magic. Add a tiny amount of platinum to hydrogen and oxygen, and they explode. Without it, they coexist peacefully. The catalyst:

  1. Lowers activation energy by providing an alternative pathway
  2. Isn't consumed in the reaction
  3. Doesn't change the overall thermodynamics (ΔG remains the same)

Enzyme catalysts in cells are masterpieces—accelerating specific reactions by factors of 10¹²-10¹⁷. They're so selective that one enzyme catalyzes one specific reaction with one specific substrate.

Socratic Questions

  1. If you measure a reaction's rate at different temperatures and find the rate increases 16-fold when temperature rises from 20°C to 40°C, what does this tell you about the activation energy?
  1. A reaction is first-order in A and second-order in B. If you increase [A] by a factor of 3 and [B] by a factor of 2, by what factor does the rate increase?
  1. A catalyst lowers a reaction's activation energy. Does it affect where equilibrium lies (does the reaction "go further")? Why or why not?
  1. Radioactive carbon-14 decays with a half-life of 5,730 years. How much of a sample remains after 11,460 years?
  1. Why does a single-step elementary reaction's order match its stoichiometry, while a multi-step mechanism's overall order might not match the balanced equation?

Solutions - Concentration effects on reaction behavior Electrochemistry - Redox reaction rates and mechanisms Chemical Equilibrium - Relationship between kinetics and thermodynamics

🃏 Flashcards — Quick Recall
Term / Concept
Rate of Reaction
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Change in concentration of reactant or product per unit time. For aA → bB: rate = −(1/a) d[A]/dt = (1/b) d[B]/dt. SI units: mol L⁻¹ s⁻¹.
Term / Concept
Rate Law
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Experimental expression rate = k[A]ˣ[B]ʸ relating reaction rate to reactant concentrations. Exponents are NOT necessarily stoichiometric coefficients — they must be determined experimentally.
Term / Concept
Order vs Molecularity
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Order = sum of powers in the experimental rate law (can be 0, fractional, or negative). Molecularity = number of species colliding in an elementary step (1, 2, or 3 only — never zero or fractional).
Term / Concept
Units of Rate Constant k
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Zero order: mol L⁻¹ s⁻¹. First order: s⁻¹. Second order: L mol⁻¹ s⁻¹. General: (mol L⁻¹)¹⁻ⁿ s⁻¹ where n is overall order.
Term / Concept
Integrated First-Order Rate Law
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ln([R]₀/[R]) = kt, or k = (2.303/t) log([R]₀/[R]). A plot of ln[R] vs t is a straight line of slope −k.
Term / Concept
Half-life of First-Order Reaction
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t₁/₂ = 0.693/k — independent of initial concentration. (For zero-order, t₁/₂ = [R]₀/2k, which DOES depend on [R]₀.)
Term / Concept
Arrhenius Equation
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k = A·e^(−Eₐ/RT). A = pre-exponential (frequency) factor, Eₐ = activation energy, R = gas constant, T = absolute temperature. Logarithmic form: ln k = ln A − Eₐ/RT.
Term / Concept
Activation Energy (Eₐ)
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Minimum extra energy reactant molecules need above their average energy to form the activated complex (transition state) and react. Higher Eₐ → slower reaction at a given T.
Term / Concept
Effect of Catalyst
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Provides an alternative pathway with lower activation energy, increasing rate. Does NOT change ΔG, ΔH, or the equilibrium constant Kc — it speeds up forward and reverse reactions equally.
Term / Concept
Pseudo First-Order Reaction
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A reaction that is genuinely higher order but appears first-order because one reactant is in large excess (its concentration stays nearly constant). Examples: acid hydrolysis of ester, inversion of cane sugar.
10 cards — click any card to flip
📝 Quick Quiz — Test Yourself
For a first-order reaction, the rate constant k = 0.0231 min⁻¹. What is the half-life?
  • A 0.693 min
  • B 30 min
  • C 0.0160 min
  • D 100 min
A reaction has rate = k[A]²[B]. By what factor does the rate change if [A] is doubled and [B] is tripled?
  • A 6 times
  • B 9 times
  • C 12 times
  • D 18 times
Which statement about a catalyst is INCORRECT?
  • A A catalyst shifts the equilibrium toward products.
  • B A catalyst lowers the activation energy.
  • C A catalyst is regenerated at the end of the reaction.
  • D A catalyst speeds up forward and reverse reactions equally.
A rate constant has units of L mol⁻¹ s⁻¹. The order of the reaction is:
  • A Zero
  • B Half
  • C First
  • D Second
Which of the following is true for an elementary reaction?
  • A Order can be fractional but molecularity cannot.
  • B Order equals molecularity.
  • C Molecularity can be zero.
  • D Order is determined from the balanced equation only.
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