1. Ch.17 Thermochemistry

2. Thermochemistry
Thermochemistry – study of the transfers of energy as heat that accompany chemical reactions and physical changes
The energy absorbed or released as heat in a chemical or physical change is measured in a calorimeter.

3. Temperature: The average kinetic energy of particles
The temperature in thermochemistry is measured in Celcius or Kelvin.
Joule – Is the SI unit for heat as well as a unit of energy

4. HEAT
Energy transferred between samples of matter because of a temperature difference.
Always moves spontaneously from higher temperature to lower temperature.

5. Specific Heat (J/g.K)
Specific Heat – energy required to raise the temperature of one gram of substance by 1oC or K.
Useful for comparing heat absorption capacities
H2O has an extreamly high specific heat (4.18J/g.K)

6. Practice
A 4gram sample of glass was heated from 274K to 314K, an increase of 40K, and was found to have absorbed 32 J of energy as heat.
What is the specific heat?
How much energy gain when heated from 314K to 344K?

7. Heat of reaction- the quantity of energy released or absorbed as heat during a chemical reaction
2 H2(g) + O2(g)  2 H2O(g) kJ
Thermochemical equation- an equation that includes the quantity of energy released or absorbed

8. The physical states of reactants and products must always be included because they influence the overall amount of energy exchanged.
What would happen to the energy of the decomposition of water if we started with ice?

9. Delta H (DH) Enthalpy-energy change-the amount of energy absorbed or lost by a system as heat during a process at constant pressure

10. DH= Hproducts - Hreactants
DH is always negative for exothermic reactions because the system loses energy
DH is always positive for endothermic reactions because the system gains energy

11. *Hess’s Law
The overall enthalpy change in a reaction is equal to the sum of enthalpy changes for the individual steps in the process.

12. *Hess’s Law
C(s) + 2H2(g) → CH4 (g) ∆Hof = ? C(s) + O2(g) → CO2(g) ∆Hof = kJ/mol H2(g) + ½ O2(g) → H2O(l) ∆Hof = kJ/mol CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) ∆Hof = kJ/mol

13. Spontaneous RxN Based on the change in energy of a reaction system
Or the randomness of the particles in a system.

14. The great majority of reactions are what? (endo- or exo- thermic?)
Do exothermic products have more or less energy then the reactants?

15. Entropy and Reaction Tendency: A tendency toward disorder
Entropy-S- a measure of the degree of randomness of the particles, such as molecules, in a system
DS is positive if there is an increase in entropy
DS is negative if there is a decrease in entropy

16. Think about a solid vs. liquid vs. gas
Which change of state has the highest DS value??

17. *Gibb’s Free Energy Equation
∆G = ∆H - T∆S -∆G = spontaneous rxn +∆G = nonspontaneous rxn

18. *Gibb’s Free Energy
For the reaction NH4Cl(s) → NH3(g) + HCl(g) , at 298K, ∆H= 176 kJ/mol and ∆S= 285 J/mol*K.
Calculate ∆G, and tell whether the reaction can proceed in the forward direction at 298K.

19. Molecular Collisions Molecules must collide in order to react.
Collision must be energetic enough.
Collision must be oriented the correct way.

20. Activation Energy
- The minimum energy required to transform the reactants into an activated complex.
Activated complex – A transitional structure that results from an effective collision and that persists while old bonds are breaking and new are forming

21. Reaction Pathway Graphs
Ea – Activation energy
Peak represents the activated complex
∆Eforward = -10kJ/mol
∆Ea = 40kJ/mol