<h2>Rate of a Chemical Reaction</h2> <p>Rate = change in concentration per unit time. For A → B: Rate = −d[A]/dt = +d[B]/dt.</p> <p><strong>Average rate</strong> vs <strong>instantaneous rate</strong> (tangent to concentration-time curve).</p> <h2>Rate Law and Rate Constant</h2> <p>Rate = k[A]ᵐ[B]ⁿ where m, n = order w.r.t. A and B (determined experimentally, not from stoichiometry).</p> <p>Units of k depend on overall order: <ul> <li>Zero order: mol L⁻¹ s⁻¹</li> <li>First order: s⁻¹</li> <li>Second order: L mol⁻¹ s⁻¹</li> </ul></p> <h2>Order vs Molecularity</h2> <table> <tr><th>Order</th><th>Molecularity</th></tr> <tr><td>Experimental</td><td>Theoretical (from mechanism)</td></tr> <tr><td>Can be fractional or zero</td><td>Always whole number (1, 2, 3)</td></tr> <tr><td>Refers to overall rate law</td><td>Refers to elementary step</td></tr> </table> <h2>Integrated Rate Laws</h2> <p><strong>Zero order:</strong> [A] = [A]₀ − kt; t½ = [A]₀/2k</p> <p><strong>First order:</strong> ln[A] = ln[A]₀ − kt; t½ = 0.693/k (independent of concentration)</p> <p><strong>Second order:</strong> 1/[A] = 1/[A]₀ + kt; t½ = 1/(k[A]₀)</p> <h2>Arrhenius Equation</h2> <p>k = A·e^(−Ea/RT)</p> <p>Taking logarithm: ln k = ln A − Ea/RT</p> <p>log(k₂/k₁) = (Ea/2.303R) × (1/T₁ − 1/T₂)</p> <p>A = frequency factor (pre-exponential factor), Ea = activation energy.</p> <h2>Effect of Temperature</h2> <p>Rate doubles/triples for every 10°C rise (thumb rule). More precisely described by Arrhenius equation. A catalyst lowers Ea without being consumed.</p>