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Reaction Energy and Reaction Kinetics

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1 Reaction Energy and Reaction Kinetics
General Chemistry Unit 12

2 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

3 Thermochemical Equations
A thermochemical equation shows the energy (enthalpy) change in the reaction Put in as reactant or product 2 H2 + O2 → 2 H2O kJ List behind as ΔH 2 H2 + O2 → 2 H2O ΔH = 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

4 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 (+,-,-)

5 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

6 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

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

8 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

9 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

10 Energy Diagrams

11 Exothermic/Endothermic

12 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

13 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

14 Factors Affecting Reaction Rate
Nature of reactants Concentration Temperature Catalysts

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

16 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

17 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

18 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

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