Reaction kinetics

🔹 Rate of Reaction

✅ 1. Key Terms:

• Rate of Reaction:
  The rate of a chemical reaction is the change in concentration of a reactant or product per unit time.
  Mathematically: Rate=change in concentration/time ​
• Frequency of Collisions:
  The number of collisions between reacting particles per unit time.
• Effective Collisions:
  Collisions that result in a chemical reaction because the particles collide with sufficient energy and proper orientation.
• Non-effective Collisions:
  Collisions where either the energy is too low or the orientation is unsuitable, so no reaction occurs.

✅ 2. Effect of Concentration and Pressure:

• Concentration (for solutions):
  Higher concentration → more particles per unit volume → increased collision frequency → more effective collisions → faster reaction.
• Pressure (for gases):
  Increasing pressure compresses the gas, increasing particle density → more collisions per second → faster rate.
• Qualitative Explanation:
  Increasing concentration or pressure increases the number of successful collisions per second, hence increases rate.

✅ 3. Calculating Rate of Reaction from Experimental Data:

• From a table of concentration vs. time: Rate=Δ[Product or Reactant] / Δt
• Units: typically mol dm⁻³ s⁻¹.
• Example: If [H₂] changes from 0.5 to 0.3 mol dm⁻³ in 20 seconds, Rate=(0.5−0.3)/20 =0.01 mol dm⁻³ s⁻¹

🔹 Effect of Temperature and Activation Energy

✅ 1. Activation Energy (E_a):

• The minimum energy that reacting particles must possess for a collision to be effective and result in a chemical reaction.

✅ 2. Boltzmann Distribution Curve:

• A graph showing the distribution of kinetic energies among particles at a given temperature.
• Features:
  • Starts at origin.
  • Peaks at most probable energy.
  • Asymptotically approaches zero.
  • Ea is marked as a threshold.
• Only particles with energy ≥ Ea can react.

Sketch characteristics:

• X-axis: Energy
• Y-axis: Number of particles
• Area under curve = total number of particles
• Shaded area after Ea = particles with enough energy to react.

✅ 3. Effect of Temperature on Rate:

• Higher Temperature:
  • Increases average kinetic energy.
  • More particles exceed Ea.
  • Curve flattens and shifts right.
  • Increases frequency of effective collisions.
• Explanation Using Boltzmann:
  • At higher T, greater area under the curve beyond Ea.
  • So, more reacting particles.

🔹 Catalysts

✅ 1. Terms:

• Catalyst:
  A substance that increases the rate of a reaction without being consumed in the process.
• Catalysis:
  The process of accelerating a chemical reaction using a catalyst.

✅ (a) Catalyst Provides an Alternative Pathway:

• Catalyst provides a different reaction mechanism with lower activation energy.
• More particles now have energy ≥ new Ea → reaction speeds up.

✅ (b) Catalyst and Boltzmann Distribution:

• Lower Ea = more particles above threshold.
• Shift Ea to left on the Boltzmann diagram → larger shaded area of effective collisions.

✅ (c) Reaction Pathway Diagram:

• Without Catalyst:
  • High activation energy peak.
• With Catalyst:
  • Lower energy peak.
• Diagram Labels:
  • Reactants
  • Products
  • Ea (uncatalyzed)
  • Ea (catalyzed)
  • ΔH (enthalpy change — same in both)

🔄 Homogeneous vs. Heterogeneous Catalysts:

• Homogeneous Catalyst:
  • Same phase as reactants (e.g., aqueous acid catalyzing esterification).
• Heterogeneous Catalyst:
  • Different phase (e.g., solid Pt catalyzing H₂ + C₂H₄).

✅ Summary Table:

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