1. Thermodynamics
Is it hot in here or what?

3. Energy Many forms and sources
Thermochemistry is interested in heat exchanges
Breaking bonds takes energy (endothermic)
Making bonds produces energy (exothermic)
SI Unit is the Joule (J), common unit calorie
4.18J = 1 cal
4.18 kJ = 1 kcal

4. Temperature and Energy Change
Every substance has a unique ability to absorb and release energy.
Called heat capacity.
Water has a very large value.
Specific Heat (cp) is the measure of how much energy a substance will absorb or release to change its temp. (4.18J/g°C for water)

5. Q = m • cp • ΔT Energy = mass x spec. heat x temp. change
Ex: If 20 g of water changes its temp from 25.0C to 30.0C, what amount of energy is absorbed?
Q = (20.0g) x (4.18 J/g C) x (5.0 C)
Q = 418 J
Practice!!!!

6. Heat of Reaction Energy change that occurs during a chemical reaction
Called enthalpy
Symbol (ΔH) change in enthalpy
Negative ΔH value = exothermic reaction
Positive ΔH value = endothermic reaction

8. Energy Reactants Overall energy change Products Reaction coordinate
Exothermic reaction
Energy
Reactants
Overall energy change
Products
Reaction coordinate

9. Endothermic reaction
Energy
Overall energy change
Reaction coordinate

10. Heat of Formation ∆Hf°
The energy change associated with the formation of one mole of a compound from its elements.
Sometimes uses fractional coefficients
Ex. H2(g) + ½ O2(g) → H2O (l)
∆Hf° = kJ/mol exothermic (comp. is stable)
Practice this skill. Important.

12. Thermochemical Equation
An equation which includes the heat energy as a reactant (endothermic) or product (exothermic)
H2(g) + ½ O2(g) → H2O (l) kJ
Notice that the energy is written as a + even though the ∆Hf° = kJ/mol

13. Calculation of ΔH reaction
Two methods – both work
Hess’ Law and a shortcut
Learn both (use the shortcut more often)

14. Hess’ Law - finding ΔH
Says that if you can write a reaction in steps that add up to the original reaction, the ΔH of the overall reaction is the sum of the ΔH of the steps. (route independent)
Guidelines
1. If the reaction is reversed, so is the sign of ΔH.
2. If coefficients are multiplied, so is ΔH.

15. Example: find ΔH CH4(g) + 2O2(g) → CO2(g) + 2H2O(l)
Solution
1. Write ΔHf eq. for each compound
C(s) + 2H2(g) → CH4(g) ΔH = kJ
C(s) + O2(g) → CO2(g) ΔH = kJ
H2(g) + ½ O2(g) → H2O(l) ΔH = kJ

16. solution for: CH4(g) + 2O2(g) → CO2(g) + 2H2O(l)
2. Multiply by any coefficients in the orig. eq.
2H2(g) + O2(g) → 2H2O(l) ΔH = kJ
3. Reverse reactions to get reactants on left
CH4(g) → C(s) + 2H2(g) ΔH = kJ
4. Add the resulting equations together……

17. Orig. eq. CH4(g) + 2O2(g) → CO2(g) + 2H2O(l)
CH4(g) → C(s) + 2H2(g) ΔH = kJ
2H2(g) + O2(g) → 2H2O(l) ΔH = kJ
C(s) + O2(g) → CO2(g) ΔH = kJ
_____________________________________
CH4(g) + 2O2(g) → CO2(g) + 2H2O(l)
(same as orig.) so add up ΔH = kJ

18. Hess’ Law
Tedious, but it works…….
Practice…….

19. ALTERNATIVE TO HESS’ LAW
Handy shortcut to use for calculations.
ΔH = Σ ΔHf (products) - Σ ΔHf (reactants)
Remember to multiply values by any coefficients in the balanced equation.
ΔH = ( kJ + 2 x kJ) – (-74.8 kJ)
ΔH = kJ

20. SPONTANEOUS REACTIONS
Proceed without outside assistance (beyond the initial Ea) -- they just happen.
Spontaneous chemical reactions generally occur if the products have lower PE than the reactants.
-sometimes this is not the case ( ex---H2O(s) only forms if the temp is less than 0 C)
-reason for this is the other driving force in nature

21. Entropy A measure of the disorder or randomness of a system.
-represented by S (change in ΔS)
ΔS = Σ Sf(products) - Σ Si(reactants)

22. All entropies (Sf and Si) are positive, but …ΔS can be negative or positive.
if ΔS is negative ===> Less disorder (lower entropy)
Example: crystallization of a liquid
if ΔS is positive ===> More disorder (more entropy)
Example: evaporation of a liquid….. favored in nature!

23. Prediction of entropy change
synthesis of compounds decreases entropy (due to bonding)
decomposition of compounds increases entropy
mixing a solute and solvent increases entropy
Solid→ Liquid → Gas increases entropy
Ie. evaporation of a liquid increases entropy

25. So>>>>>>
Two forces influence the direction of a spontaneous reaction: ΔH and ΔS
when ΔH decreases; lower energy results.
when ΔS increases; higher disorder results

26. These two factors compete
Like a tug of war.
Usually the enthalpy prevails.

27. Gibbs (Free) Energy ΔG
An equation called the “Gibbs equation” compares the values of ΔH and Δ S.
ΔG = ΔH - T ΔS
if ΔG is negative; the reaction is spontaneous
if ΔG is positive: the reaction is not spontaneous, and will not occur.

28. 4 Possible Cases ΔH ΔS Result: - + reaction favored + - no reaction
rx if ΔH is large
rx if ΔS is large

29. Alternative to the Gibbs equation:
ΔG = ΣΔGf (products) - ΣΔGf (reactants)

30. Collision Theory
In order to react molecules and atoms must touch each other.
They must hit each other hard enough to react.
Anything that increase these things will make the reaction faster.

31. Things that Effect Rate
TEMPERATURE
Higher temperature = faster particles.
More and harder collisions.
Faster Reactions.
CONCENTRATION
More concentrated = molecules closer together.
Collide more often.
Faster reaction.

32. Things that Effect Rate
PARTICLE SIZE
Molecules can only collide at the surface.
Smaller particles = bigger surface area.
Smaller particles = faster reaction.
Smallest possible is atoms or ions.
Dissolving speeds up reactions.
Getting two solids to react with each other is slow.

33. Things that Effect Rate
CATALYSTS- substances that speed up a reaction without being used up.(enzyme).
Speeds up reaction by giving the reaction a new path.
The new path has a lower activation energy.
More molecules have this energy.
The reaction goes faster.
Inhibitor- a substance that blocks a catalyst.