1. Introduction to Thermochemistry

2. Law of Conservation of Energy
In any chemical or physical process, energy can not be created nor destroyed.
All the energy in a process can be accounted for as work, light, stored energy, or heat.

3. Temperature
Temperature is a measurement of the average speed that molecules are moving.
It’s an average b/c not all molecules are moving at the same speed at the same time.
Measured in Kelvin (S.I) or degrees Celsius
Kelvin= °C ** Memorize This**
If room temp is 20°C, what is it in Kelvin?
If absolute zero is 0 Kelvin, what is in Celsius?

4. Heat
The energy that is transferred from one body to another because of temperature differences
Flows from warmer → cooler object
Only changes caused by heat can be detected – like changes in temperature

5. Units of Heat Joule (J) is the SI unit of heat and energy.
calorie (cal) is another commonly used unit of energy defined as the quantity of heat needed to raise the temperature of 1g of water 1oC.
4.184 J = 1 cal
fyi, the calories in food are kilocalories!!

6. How many joules are in 250 Calories?

7. Enthalpy ∆H
Enthalpy (H) is a measure of the internal energy of a system plus temperature and pressure.
We can only measure changes in enthalpy (∆H)
Heat content (q) is the same as enthalpy for systems at a constant pressure
Therefore q and H are often used interchangeably.
q = H at constant pressure

8. System vs. Surroundings Which way does the heat flow?
In any chemical reaction or change of physical state, heat is either released or absorbed.
In studying which direction heat flows, we use these two definitions:
the system - the part of the universe on which you focus your attention
the surroundings - includes everything else in the universe

9. Exothermic or Endothermic
Endothermic Process: Heat Absorbed
Heat going into the system from the surroundings.
q or +∆H is positive
Ex: Can be chemical (ex: Ice pack) or physical (melting)
Exothermic Process: Heat Released
Heat leaving the system and going into the surroundings.
q or -∆H is negative
Can be chemical (ex: combustion) or physical (freezing)

10. Thermochemical Equations
A thermochemical equations shows the heat change.
The heat is usually given in kJ and can be included in the reaction itself or directly after the equation.
Can be used for both phase changes and chemical changes
The heat change for a reaction is called Hrxn.

11. Enthalpies of phase changes
Hfus is the amount of heat required to melt (fuse) one mole of a substance
Hvap is the amount of heat required to vaporize one mole of a substance

12. +Hfusion = - ∆Hsolidification +Hvaporization = - ∆Hcondensation
Molar Relationships
Since freezing and melting both occur at the same temperature Hfusion is the same but opposite as the ∆Hsolidification
+Hfusion = - ∆Hsolidification
Same is true for vaporization & condensation
+Hvaporization = - ∆Hcondensation

13. Heat of Physical Changes Table
Substance
Melting point °C
Heat of fusion (kJ/kg)
Heat of solidification
(kJ/kg)
Boiling point °C
Heat of vaporization (kJ/kg)
Heat of condensation
Helium
5.23
-5.23
20.9
-20.9
Hydrogen
58.6
-58.6
452
-452
Nitrogen
25.5
-25.5
201
-201
Oxygen
13.8
-13.8
213
-213
Ethyl alcohol
-114
104.2
-104.2 78
854
-854
Mercury
-39
11.8
-11.8
357
272
-272
Water
0.00
334
-334
100.00
2256
-2256
Sulfur
119
38.1
-38.1
444.60
326
-326
Lead
327.3
24.5
-24.5
1750
871
-871
Antimony
630.50
165
-165
1440
561
-561
Silver
960.80
88.3
-88.3
2193
2336
-2336
Gold
64.5
-64.5
2660
1578
-1578
Copper
1083
134
-134
2567
5069
-5069

14. Practice Problem #1
How much heat is absorbed when 24.8 g water is evaporated? (∆Hvap = 6.01kJ/mol)

15. Practice Problem #2
If 300 kJ of heat is available, how many grams copper can be melted?(∆Hfus= 8.52kJ/mol)

16. Practice Problem #3
How much heat is released when 100 grams of ethyl alcohol condenses? (∆Hvap = 854 kJ/kg)

17. Chemical Exothermic Reactions
Energy is written as a product, or
H is negative
CaO (s) + H2O(l)  Ca(OH) kJ or
CaO (s) + H2O(l)  Ca(OH)2 H = kJ

18. Chemical Endothermic Reactions
Energy is written as a reactant
+H is positive
2NaHCO kJ  Na2CO3 + H2O +CO2 or
2NaHCO3  Na2CO3 + H2O(g) +CO2 H = 129 kJ

19. Exothermic or Endothermic?
C2H5OH(l) + 3O2(g)  2CO2(g) + 3H2O(l) kJ
PCl5(s)  PCl3(g) + Cl2 H = 87.9 kJ
H2(g) + F2(g)  2HF(g) H = -536 J
CO2(g) kJ  C(s, graphite) + O2 (g)
Exo b/c energy(enthalpy) is a product
Endo b/c energy(enthalpy) is positive
Exo b/c energy(enthalpy) is negative
Endo b/c energy(enthalpy) is a reactant

20. 2NaHCO3 (s)  Na2CO3 (s) + H2O (g) + CO2 (g) H = +129 kJ
Practice Problem #4
Calculate the kilojoules of heat required to decompose 2.24 mol sodium bicarbonate (aka baking soda)
2NaHCO3 (s)  Na2CO3 (s) + H2O (g) + CO2 (g) H = +129 kJ

22. C2H4 (g) + 3O2 (g)  2CO2 (g) + 2 H2O (l) + 1411kJ.
Practice Problem #5
How much heat is released when 8.0 g of oxygen react in:
C2H4 (g) + 3O2 (g)  2CO2 (g) + 2 H2O (l) kJ.

24. Si (s) + 2Cl2 (g)  SiCl2 (l) + 687 kJ
Practice Problem #6
Si (s) + 2Cl2 (g)  SiCl2 (l) kJ
How much heat is produced when 106 grams of chlorine reacts with silicon?

26. Potential Energy/ Reaction Pathway Diagrams
Activation energy (Ea) : energy needed to start a reaction book clip Activated Complex: the “top of the hill” reactants’ atoms start rearranging due to collisions. H: Enthalpy: change in heat H =products – reactants Catalyst: speeds up a rxn, w/o changing the reaction by lowering the activation energy!!!

28. Potential Energy Diagrams
Reactants have less energy than Products, so it had to gain energy)
Reactants have more energy than products, so it had to lose energy

29. Heating Curves
A heating curve is a graph that shows temperature changes in a substance as energy (heat) is added.
It is a plot of temperature vs. time, showing phase changes.

30. Heating Curves Slanted lines show a change in kinetic energy.
Temperature is increasing. (Kinetic energy is directly related to temperature.)
Flat lines indicate phase changes.
Temperature is NOT changing.
Energy is still being added—potential energy.
The energy is being used to break intermolecular forces.
The stronger the IMFs, the more energy needs to be added in order to complete a phase change.

31. Heating Curves a → b solid warming b → c solid becoming a liquid
Hvap
Hfus
Potential energy change
Kinetic energy change
a → b solid warming
b → c solid becoming a liquid
c → d liquid warming
d → e liquid becoming a gas
e → f gas warming

32. Drawing Heating Curves
Draw the heating curve for water showing the energy changes from liquid at 30 ºC to steam at 110 ºC.
Step 1: Draw a graph / label low and high temp from problem
Step 2: Determine any phase changes
(For water: 0 ºC and 100 ºC)
Step 3: Draw slanted lines to and from phase changes and flat lines for phase changes