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