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Reaction Energy and Reaction Kinetics General Chemistry Unit 12.

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Presentation on theme: "Reaction Energy and Reaction Kinetics General Chemistry Unit 12."— Presentation transcript:

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 H 2 + O 2 2 H 2 O kJ List behind as ΔH 2 H 2 + O 2 2 H 2 O 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: NH 4 Cl (s) NH 3(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 ΔGSpontaneous? 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 1.Calculate the ΔH. 20 kJ – 40 kJ = -20 kJ 2.Calculate the E a. 100 kJ – 40 kJ = 60 kJ 3.Calculate the E a. 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 Animation Heterogeneous – not in same phase as reactants, provides surface to give more effective collisions Catalytic ConverterCatalytic Converter Homogeneous – In same phase as reactants, makes different activated complex, returns to original form at end of reaction Demo: CatalystsCatalysts

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