Overview
This lecture introduces advanced kinetics, focusing on measuring reaction rates, interpreting rate equations, determining reaction orders, and understanding rate-determining steps using both continuous and initial rate experimental methods.
Measuring Rate of Reaction
- Rate of reaction is defined as the change in concentration of product or reactant per unit time.
- Standard unit for rate is mol dm⁻³ s⁻¹ (moles per cubic decimeter per second).
- Six main methods for measuring rates: gas volume, mass change, color intensity (colorimetry), titration, pH change, and electrical conductivity.
- Choice of method depends on the nature and precision required for the reaction.
Rate Equations and Rate Constants
- Rate equation expresses rate as a function of reactant concentrations: rate = k[A]^m[B]^n.
- 'k' is the rate constant; its unit depends on the overall order of the reaction.
- First-order: rate ∝ [A]; Second-order: rate ∝ [A]^2; Zero-order: rate is independent of [A].
- The unit of the rate constant varies (e.g., for second order: dm³ mol⁻¹ s⁻¹).
Reaction Order and Determination
- Order indicates how the rate changes as concentrations of reactants change.
- Overall order is the sum of the individual orders with respect to each reactant.
- Orders are determined experimentally, not from the stoichiometric equation.
- Zero order: no effect of concentration; First order: rate doubles if concentration doubles; Second order: rate quadruples if concentration doubles.
Rate Determining Step (RDS)
- The slowest step in a reaction mechanism controls the overall reaction rate.
- Only reactants in the RDS appear in the rate equation; zero-order reactants do not.
- Mechanistic steps are inferred from experimental rate data.
Experimental Methods for Determining Order
Continuous Monitoring
- Take samples from a single reaction over time, quench, and analyze (e.g., by titration).
- Plot concentration vs. time; constant half-life signifies first-order, doubling half-life indicates second-order, straight line indicates zero-order.
- Draw tangents to curves to find instantaneous rates.
Initial Rate Method
- Repeat the reaction several times, varying the concentration of one reactant each time.
- Compare changes in rate to changes in concentration to deduce orders.
- Linear relationship (rate doubles with concentration) indicates first order; quadruple increase indicates second order.
Worked Example: NO + H₂ Reaction
- Experimentally compare runs with varying NO and H₂ concentrations to determine orders: NO is second order (rate increases by 16 when [NO] increases by 4), H₂ is first order (consistent with observed rate changes).
- Write rate equation: rate = k[NO]²[H₂].
- Calculate k using experimental data and assign correct units based on overall reaction order.
Key Terms & Definitions
- Rate of Reaction — change in concentration of a reactant or product per unit time.
- Rate Equation — mathematical relationship linking rate to reactant concentrations.
- Rate Constant (k) — proportionality constant in the rate equation.
- Order of Reaction — power to which the concentration of a reactant is raised in the rate equation.
- Half-life — time taken for concentration of a reactant to halve.
- Rate Determining Step (RDS) — the slowest step in a reaction mechanism controlling the overall rate.
Action Items / Next Steps
- Memorize the units for rate constants for different reaction orders.
- Practice deriving rate constant units from rate equations.
- Review and analyze sample data sets using both the continuous and initial rate method.
- Complete core practical 9a and 9b as discussed.