Reaction Energy and Reaction Kinetics General Chemistry Unit 12

Driving Forces Enthalpy and Entropy Enthalpy (heat of reaction) is the amount of energy released or absorbed during a chemical reaction Symbol is ΔH Think of it as energy needed

Thermochemical Equations A thermochemical equation shows the energy (enthalpy) change in the reaction Put in as reactant or product 2 H2 + O2 → 2 H2O + 483.6 kJ List behind as ΔH 2 H2 + O2 → 2 H2O ΔH = -483.6 kJ If energy is released (product) the reaction is exothermic and ΔH is negative If energy is absorbed (reactant) the reaction is endothermic and ΔH is positive

Entropy Entropy is a measure of randomness, tendency toward disorder Symbol is ΔS More disorder = more entropy If reaction leads to more disorder, the entropy change (ΔS) is positive, if it becomes more ordered, ΔS is negative Example: melting ice, condensing water, cleaning your room (+,-,-)

Free Energy (ΔG) Free energy combines enthalpy and entropy to measure the spontanaeity of a reaction Gibbs Free Energy Equation: ΔG = ΔH - T ΔS (T is in Kelvin: +273 to ºC) If ΔG is negative, reaction is spontaneous If ΔG is positive, reaction is NOT spontaneous, but would be spontaneous in the reverse direction

Example Find ΔG for the reaction: NH4Cl(s) → NH3(g) + HCl(g) Using the following data: ΔH = 176 kJ, ΔS = 285 J/K, T = 25ºC Solution: (Change to kJ and K) ΔG = 176 kJ – (298 K)(.285 kJ/K) ΔG = 176 kJ – 84.9 kJ = 91 kJ NOT spontaneous

Comparison of Signs ΔH ΔS ΔG Spontaneous? - + - ALWAYS spont. + - + NEVER spont. - - - / + Spont. at low T + + - / + Spont. at high T

Reaction Mechanisms Step-by-step sequence that occurs to create the products Intermediates may form that do not appear in overall reaction – they are used up in another step Homogeneous reaction: all reactants in same phase Heterogeneous reaction: reactants in different phases Rate-determining step: slowest step of reaction mechanism

Activation Energy Minimum energy to make the reaction go (form activated complex which allows reaction to proceed) Reaction needs: Enough energy Proper orientation of molecules – must hit each other at correct spot

Energy Diagrams

Exothermic/Endothermic

Energy Example Calculate the ΔH. Calculate the Ea. Calculate the Ea‘. 20 kJ – 40 kJ = -20 kJ Calculate the Ea. 100 kJ – 40 kJ = 60 kJ Calculate the Ea‘. 100 kJ – 20 kJ = 80 kJ

Reaction Rate Rate can be defined in terms of molar concentration (M) for the disappearance of a reactant or the appearance of a product Concentration shown as: [HCl] = 0.1 means the molar concentration of HCl is 0.1 M

Factors Affecting Reaction Rate Nature of reactants Concentration Temperature Catalysts

Nature of Reactants Ionic – almost instantaneous Molecular – slower (bonds must break and reform) Surface area – rate increases with greater surface area

Concentration Measured in molarity [A] Increasing the concentration of reactants increases the rate Rate law: Rate = k[A]m[B]n The exponents m and n must be determined experimentally

Temperature Increasing the temperature gives more collisions between molecules This leads to the formation of more activated complexes and this causes the rate to increase ↑ T → ↑ collisions → ↑ complexes → ↑ rate

Catalysts Catalyst – increase reaction rate without being used up Lower the activation energy Animation Heterogeneous – not in same phase as reactants, provides surface to give more effective collisions Catalytic Converter Homogeneous – In same phase as reactants, makes different activated complex, returns to original form at end of reaction Demo: Catalysts