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Reactivity 2. How much, how fast and how far?

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  1. 2.1 Amount of chemical change

    1. • Show ratio of reactants and products • Deduce balanced equations, including state symbols, from named reactants and products

    2. • Determine masses, volumes, concentrations • Use balanced equations for reacting quantity calculations

    3. • Determines theoretical yield • Identify limiting and excess reactants from given data

    4. • (Experimental yield / Theoretical yield) × 100% • Solve problems involving theoretical, experimental, and percentage yields

    5. • Measure of green chemistry efficiency • Calculate atom economy and link low atom economy to waste

  2. 2.2 Rate of chemical change

    1. 2.2.1Rate of reaction

      • Change in concentration per unit time • Determine rates from tangents of concentration, volume, or mass graphs

    2. 2.2.2Collision theory

      • Sufficient energy + proper orientation required • Kinetic energy and temperature (Kelvin) • Explain successful collisions using activation energy and orientation

    3. 2.2.3Factors affecting rate

      • Pressure, concentration, surface area, temperature, catalyst • Predict and explain rate changes when conditions change

    4. 2.2.4Activation energy (Ea)

      • Minimum energy for successful collision • Maxwell-Boltzmann distribution curves • Use Maxwell-Boltzmann curves to show temperature effects on successful collisions

    5. 2.2.5Catalysts

      • Provide alternative pathway with lower Ea • Energy profiles with/without catalysts • Explain catalyst effects on energy profiles and Maxwell-Boltzmann distributions

    6. 2.2.6(HL)—Reaction mechanisms

      • Series of elementary steps • Slowest step = rate-determining • Intermediates vs. transition states • Evaluate proposed mechanisms against kinetic and stoichiometric data

    7. 2.2.7(HL)—Energy profiles

      • Show Ea and transition state of rate-determining step • Construct and interpret multistep energy profiles from kinetic data

    8. 2.2.8(HL)—Molecularity

      • Number of particles in elementary step • Unimolecular, bimolecular, termolecular

    9. 2.2.9(HL)—Rate equations

      • Determined experimentally, depend on mechanism • Deduce rate equations from experimental data

    10. 2.2.10(HL)—Reaction order

      • Exponent in rate equation • Zero, first, second order • Overall order = sum of individual orders • Analyse concentration-time and rate-concentration graphs

    11. 2.2.11(HL)—Rate constant (k)

      • Temperature dependent • Units from overall order • Solve rate equation problems, including units of k

    12. 2.2.12(HL)—Arrhenius equation

      • Temperature dependence of k determines Ea • Linear form: ln k vs. 1/T

    13. 2.2.13(HL)—Arrhenius factor (A)

      • Frequency of proper collision orientations • Determine activation energy and Arrhenius factor from data

  3. 2.3 Extent of chemical change

    1. • Forward rate = backward rate (closed system) • Describe physical and chemical systems at equilibrium

    2. • K determined from stoichiometry • Deduce K expressions for homogeneous reactions

    3. • Magnitude indicates extent of reaction • Temperature dependent • Kreverse = 1/Kforward • Interpret K << 1, K < 1, K = 1, K > 1, and K >> 1

    4. • Predict effects of changes in concentration, temperature, pressure • Include effects on K and equilibrium composition

    5. • Non-equilibrium concentrations • Compare Q to K to predict direction • Calculate Q and predict direction toward equilibrium

    6. • K with initial and equilibrium concentrations • Use small-K approximations; quadratic equations are not expected

    7. • Both measure equilibrium position • Calculate ΔG⦵ from ΔG⦵ = −RT ln K