1. Chapters 16 & 17
Thermochemistry

2. Thermochemistry The study of energy transfers as heat (Q) Heat (Q)
-The energy transferred when there is a difference in temperature between two areas
-Measured in Joules (J) or kilojoules (kJ)

3. Temperature Measure of AVERAGE kinetic energy of particles
Faster particles = higher temps
Slower particles = lower temps
Measured in ℃elsius or Kelvins.
T(°C) = (T(°F) - 32) × 5/9
K = ℃

4. Heat of Reaction Some reactions release or absorb heat Exothermic
-Heat energy is released (product)
-Products lower in energy than reactants
-Feels warm
Endothermic
-Heat energy is absorbed (reactant)
-Products higher in energy than reactants
-Feels cold

5. How do we include heat in a reaction?
1. Show it as a reactant or product: 2H2(g) + O2(g)  2H2O(g) kJ 2H2O(g) kJ  2H2(g) + O2(g) *This can be used as a conversion factor! (2 moles H2O = kJ heat)
Exothermic
Endothermic

6. How do we include heat in a reaction?
2. Write it separately as a H value (enthalpy change) Enthalpy Change, H H = Hproducts – Hreactants 2H2(g) + O2(g)  2H2O(g) H = kJ/mol 2H2O(g)  2H2(g) + O2(g) H = kJ/mol
H = exothermic
+H = endothermic
Exothermic
Endothermic

7. How do we show heat in a reaction?
3. Energy diagrams – show H visually

8. Energy Diagram
Activated complex
-Transition state
-Unstable

9. Another energy diagram
Transition states
Intermediate
Energy
Reactant
Product
Reaction Progress

10. Specific Heat (cp) Q Cp = Q = Cpmt mt heat change in temperature
Quantity of heat needed to raise the temp of 1 gram of a substance by 1C or 1K
Units = J/gC or J/gK
Different substances have different specific heats
Water – high specific heat; Iron – low specific heat Q
Cp =
Q = Cpmt
mt
heat
change in temperature
mass

11. Specific Heat (J/g℃) Water (l) 4.18 Water (s) 2.06 Aluminum 0.897
Steel
0.490
Iron
0.449
Zinc
0.388
Copper
0.385
Brass
0.380
Tin
0.228
Lead
0.129

13. Calorimeter – reaction chamber
Used to measure the energy absorbed or released as heat in a chemical or physical change.
Usually substances are immersed in a known quantity of water.
Energy given off by the reaction is absorbed by the water, and the Δt of the water is measured.
From the Δt of the water, it is possible to calculate the Q of the reaction.

14. Combustion Calorimeter

15. Examples: pg.) 534
Determine the specific heat of a material if a 35g sample absorbed 96J as it was heated from 293K to 313K.

16. If 980kJ of energy are added to 6
If 980kJ of energy are added to 6.2L of water at 291K, what will the final temperature of the water be?

17. Calculating ΔH of Reaction
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.
Ex.) Formation of methane gas:
CH4(g) + 2O2(g)  CO2(g) + 2H2O(l) ΔH = kJ
C(s) + O2(g)  CO2(g) ΔH = kJ
2H2(g) + O2(g)  2H2O(l) ΔH = kJ
C(s) + 2H2(g)  CH4(g) ΔH = ?

18. Enthalpy Formation C + O2  CO2 ΔH = -393.5 kJ
H2+ O2 H2O ΔH = kJ
C5H12+ 8O2 5CO2 + 6H2O ΔH = kJ
C(s) + H2(g)  C5H12(g) ΔH = ? kJ

19. Combustion of C8H18 C(s) + O2(g)  CO2(g) ΔH = -393.5 kJ
H2(g) + 0.5O2(g)  H2O(l) ΔH = kJ
8C(s) + 9H2(g)  C8H18(l) ΔH = kJ

20. Driving Forces of Reactions
Not all reactions are spontaneous
Spontaneity depends on
1. Enthalpy (H)
2. Entropy (S)

21. 1. Enthalpy (H) -H = spontaneous +H = not spontaneous
Lower energy is more favorable
(So, exothermic reactions usually “work”)
Exothermic
-H
Heat released (is a product)
Feels hot
-H = spontaneous
+H = not spontaneous

22. 2. Entropy (S) -S = not spontaneous +S = spontaneous
Entropy is a measure of DISORDER in a system
Units of J/molK
T = S
gas = S
Dissolving = S
More disorder = more favorable/spontaneous
2NH4NO3(s)  2N2 (g) + 4H2O (l) + O2(g)
-S = not spontaneous
+S = spontaneous
Is this spontaneous?
Yes! Increasing disorder!!

23. What if H and  S compete?
Put them together!!
Use Gibb’s Free Energy calculation

24. Gibb’s Free Energy G -G = spontaneous +G = not spontaneous
Combines H and S to determine if a reaction will be spontaneous
G = H - T S
-G = spontaneous
+G = not spontaneous

25. Collision Theory Collision Theory Activation energy, Ea :
Particles must collide with enough energy and force at a proper orientation in order for a reaction to occur
Activation energy, Ea :
- Energy required for a reaction to occur

27. Reaction Rate Influenced by 5 factors: Nature of the reactants
Bond strength, size of molecules, etc.
Surface Area
surface area, rate
Temperature
T, KE, collisions, rate
Concentration
[ ], rate
Catalyst – lowers activation energy

28. 5. Catalyst
- Changes rate but is NOT part of or used up in a reaction

29. Chapter Review
Pg 552
#’s 1, 3-4, 7-10, 12, 23-25
Due day of test

30. Keq – Equilibrium Constant
Gives a picture of what side of an equilibrium reaction is favored
H2 + I2 <--> 2HI
[products]
Keq =
[reactants]
Equilibrium expression
Keq > 1 = Products favored
Keq < 1 = Reactants favored
[HI]2
Keq =
[H2] X [I2]

31. LeChatelier’s Principle
If a system at equilibrium is disturbed, the equilibrium will shift to relieve the stress
PCl5 (g) + heat PCl3 (g) + Cl2 (g)
What happens if:
Cl2?
heat?
heat?
PCl3?
Add a catalyst?

32. Reversible reaction
Products can react to form reactants

33. Equilibrium Rate of forward reaction = rate of reverse reaction
[products] and [reactants] is unchanged
Is dynamic (always changing)
Sometimes favors one side:
HBr + H2O H3O+ + Br-
H2CO3 + H2O H3O+ + HCO3-