1. Unit 13 Thermochemistry

2. Energy The ability to do work or cause a change
Often measured in joules (J)
Law of Conservation of Energy – energy is neither created nor destroyed
It is often transferred & can change form

3. Enthalpy (H)
Total energy content in a sample under a constant pressure
The change in enthalpy (H) is often used interchangeably with heat (q)
Often measured in joules

5. Units of Energy Energy can also be measured in calories
calorie is amount of heat to change 1 g of water by 1 C
Food Calories are kilocalories
1 Calorie = 1 kilocalorie = 1000 calories
1 cal = J
1 Cal = 4184 J

6. Heat in phase changes
Remember, the temperature does not change during a phase change

7. Latent heat
Heat absorbed or released by a substance during a process that does not involve a change in temperature
Solid  Liquid = Latent Heat of fusion (Lf)
Liquid  Gas = Latent Heat of vaporization (Lv)
Varies for substances based on IMF & mass

9. Calculating latent heat
Q=m Lf OR Q=m Lv
How much energy does it take to melt 16 g of ice at it’s melting point? The Lf of ice is 80 cal/g
Q = m Lf
Q = 16g * 80 cal/g
Q = 1280 calories OR Calories

10. Molar heat of fusion ( H fus)
Heat absorbed by 1 mole of a solid as it melts to a liquid at a constant temperature
Molar heat of solidification ( H solid)
Heat lost when 1 mole of a liquid becomes a solid
( H fus) = ( H solid) but opposite sign

12. Using molar heat of fusion
How many grams of ice at 0 deg C melts if 2.25 kJ heat added? The molar heat of fusion of water is 6.01 kJ/mol.
2.25 kJ g
mol 1
6.01
mol kJ 18 1

13. Molar heat of vaporization
Amount of heat to vaporize 1 mole of a substance
Equal but opposite to the molar heat of condensation

14. UNDERSTANDING CHECK
The heat of vaporization of water is kJ/mol. How much heat is required to turn 14 g of water into vapor at the boiling point?

15. Heat (q)
Heat – energy transfer between objects at different temperatures
Goes from warmer object to cooler object and not vice-versa
Temperature (T) – measure of how hot or cold OR measure of average kinetic energy
Higher temp = faster movement

16. Heat vs. Temperature Heat is an EXTENSIVE PROPERTY
Physical property that depends on the amount of sample present
Match vs campfire
Temperature is an INTENSIVE PROPERTY
Physical property that does not depend on the amount of sample

17. Exothermic Reaction Surroundings System Energy
Change in matter where E is released, often as heat
Surroundings get hotter
q or H of the system is negative
System
Surroundings
Energy

18. Exothermic Reactions
Joules of energy created are written with the products of the balanced reaction
On a graph, the products are lower in energy than the reactants
Makes sense because energy is RELEASED during the reaction

19. C + O2 ® CO2
+ 395 kJ
Energy
Reactants
Products
C + O2
-395kJ
CO2

20. Endothermic Surroundings System Energy
Change in matter in which energy is absorbed, often as heat
Surroundings get cooler
q or H of the system is positive
System
Surroundings
Energy

21. Endothermic
Joules of energy absorbed are written with the reactants of the balanced reaction
On a graph, the products are higher in energy than the reactants
Makes sense because energy is ABSORBED during the reaction

22. CaCO3 + 176 kJ ® CaO + CO2 CaCO3 ® CaO + CO2 Energy Reactants Products

24. Heat capacity (c) How much heat it takes to heat an object by 1C
Affected by identity of the substance (sand vs. water) & amt present (ocean vs. pan)

25. Heat capacity
Specific heat capacity (c) is the amount needed to heat 1 g of a substance by 1C
Affected by identity of the substance
The higher the specific heat, the more energy it takes to change its temperature
Formula q = m c T
For water, c = 1 cal/gºC or J/gºC

26. How much heat is needed to change the temperature of 12 g of silver with a specific heat of cal/gC from 25C to 83 C?
Q = m c T
m = 12 g
c = cal/gC
T = 58 C (no need to change to Kelvin)

27. If you put 6500 J of heat into a 15 g piece of Al at 25 C , what will the final temperature be? ( c = 0.90 J/gC )
Q = m c T
m = 15 g
c = 0.90 J/gC
Q = 6500 J

28. Applying specific heat
Equal masses of copper and water are heated until the temperature is increased by 10 degrees. Why did the copper’s temperature increase more quickly?
Water has a very HIGH heat capacity and therefore changes temp SLOWLY
Walking on water or sand on a hot day?

29. 1st law of thermodynamics
Law of Conservation of Energy – The change in the internal energy of a system is equal to the amount of energy added by heating minus the amount lost by the work done by the system
Eating a cheeseburger gives you the amount of energy equal to what’s in the burger minus the energy used to chew & digest

30. Calorimetry Measures heat using a device called a calorimeter
Measure temp before and after reaction
Since the change in enthalpy (H) can be used interchangeably with heat (q), we can use the equation q = m c T
Water
c = 1 calorie/gram * degree Celsius
c = Joule/gram * degree Celsius

31. Bomb calorimeter

32. Coffee Cup Calorimeter

33. Example
A chemical reaction is carried out in a coffee cup calorimeter. There are 75.8 g of water in the cup, and the temperature rises from 16.8 ºC to 34.3 ºC. How much heat was released?