Chemical Kinetics

Main Ideas / Lessons Conveyed

• Chemical kinetics is the branch of chemistry that studies:
  • Rate of chemical reactions (how fast reactions happen)
  • Factors that affect reaction speed, including:
    • Concentration
    • Temperature
    • Catalyst
    • Surface area
    • Pressure (mainly for gases)
    • Nature of reactants
  • Why reactions occur at a given speed, linked to:
    • Molecular collisions
    • Energy barriers

Rate of Reaction (Definition)

The rate of reaction is the change in concentration of reactants and/or products per unit time.

For concentration terms:

• Reactant concentration decreases → rate expression uses a negative sign (disappearance)
• Product concentration increases → rate expression uses a positive sign (appearance)

Mathematical Form

Conceptually:

• Rate ∝ change in concentration / change in time

Typically expressed as: [ \text{Rate} = \frac{\Delta[\text{concentration}]}{\Delta t} ]

Concentration units mentioned:

• moles per liter
• moles per dm³ (dm³)

Rate Equation / Rate Law

A rate equation (rate law) links reaction rate to the concentration(s) of reactant(s): [ R = k[\text{A}]^{m}[\text{B}]^{n} ] (conceptually stated)

• The exponents (m, n, …) represent the order of reaction with respect to each reactant.
• The overall order is the sum of the individual orders.

Rate constant (k):

• Changes only with temperature (as stated)
• Does not change with concentration (in the taught context)

Orders of Reaction (Examples)

• Zero order

  • Rate is independent of concentration
  • [ R = k ]

• First order

  • Rate is proportional to concentration
  • [ R = k[A] ]

• Second order (examples)

  • [ R = k[A]^2 ]

  • or [ R = k[A][B] ] depending on which reactant terms appear

• Fractional order

  • Example discussed: order like (1/2)
• Third order (example discussed)
  • [ R = k[A]^3 ]

Experimental Determination: Initial Rate Method

To find reaction order:

• Measure the initial rate at different initial concentrations.
• Change one initial concentration at a time while keeping others constant.
• Observe how the initial rate changes:
  • If concentration doubles and rate doubles → first order
  • If concentration doubles and rate becomes 4× → second order
  • If concentration doubles and rate becomes 8× → third order
  • If rate stays constant when concentration changes → zero order

A concentration vs. rate table is used to organize results and repeat for different initial values.

Graphical Behavior (Qualitative)

• Concentration vs time graphs can be used to infer order behavior.
• Half-life trends emphasized:
  • Zero order: half-life decreases as concentration decreases (as stated)
  • First order: half-life treated as constant (as stated)
  • Second order: half-life increases as concentration decreases (as stated)
  • Fractional order: graphs described as non-standard / not simple straight lines

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Factors Affecting Reaction Rate

Concentration

• Higher reactant concentration → more particles in the same volume → more effective collisions → rate increases
• Lower reactant concentration → fewer collisions → rate decreases

Temperature

• Higher temperature → particles have higher kinetic energy → more collisions
• More collisions exceed activation energy → reaction proceeds faster
• Lower temperature → slower reaction

Catalyst / Catalysis

• A catalyst increases reaction rate without being consumed
• Provides an alternate pathway with lower activation energy

Surface Area (especially for solids)

• Larger surface area → more exposed particles → more collisions → rate increases
• Smaller surface area → rate decreases

Pressure (mainly gases)

• Increasing pressure increases gas particle concentration → more collisions → rate increases

Nature of Reactants

• Some reactants react faster by nature than others
• Example given:
  • Ionic compounds react faster than covalent compounds (as taught)

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Methodology: Determining Order (Initial Rate Method)

1. Prepare the reaction under controlled conditions.
2. Measure and record:
   • Initial concentration(s) of reactants
   • Initial rate (initial reaction rate)
3. Change only one initial concentration at a time; keep others constant.
4. Compare how initial rate changes when concentration doubles, triples, etc.
5. Infer order from the concentration–rate relationship:
   • Zero order: rate unchanged with concentration changes
   • First order: rate changes proportionally with concentration
   • Second order: rate changes as the square (e.g., doubling concentration → 4× rate)
   • Third order: rate changes as the cube (e.g., doubling concentration → 8× rate)
6. Use a concentration–rate table to organize data and calculate order.

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Rate Equation Formation (Conceptual)

• Write the general rate law form: [ R = k[A]^m[B]^n ]

• Determine exponents (m, n) experimentally (from initial rate behavior).

• Compute overall order as: [ m + n + \dots ]

• Use the rate constant (k) for the reaction at the given temperature.

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Activation Energy and Energy Diagram

Activation Energy ((E_a))

• Defined as the minimum energy required for reactant molecules to form the activated complex
• Without reaching (E_a), the reaction cannot proceed

Activated Complex / Transition State (Concept)

• A temporary, short-lived, unstable arrangement of atoms
• Occurs at the highest point on the energy barrier
• Can:
  • form products, or
  • revert back to reactants

Energy Barrier Idea

• Reactants → rise to the activated complex (peak) → fall to products
• Catalysts lower activation energy, enabling faster reaction (as stated), without fundamentally changing equilibrium.

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Arrhenius Equation (As Mentioned)

• The rate constant (k) depends on temperature.

• Arrhenius relation given: [ k = A e^{-E_a/(RT)} ]

• Intended relationship emphasized:

  • Higher temperature → larger (k) → faster reaction

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