1. Thermochemistry

2. Thermochemistry
Thermochemistry is the study of heat changes that accompany chemical reactions and phase changes.
In thermochemistry, the system is the specific part of the universe that contains the reaction or process you wish to study.

3. universe = system + surroundings
Thermochemistry
Everything in the universe other than the system is considered the surroundings.
Therefore, the universe is defined as the system plus the surroundings.
universe = system + surroundings

4. 1st Law of Thermodynamics
1st Law of Thermodynamics – energy is conserved
The law of conservation of energy states that in any chemical reaction or physical process, energy can be converted from one form to another, but it is neither created nor destroyed.

5. Energy Energy – the ability to do work or produce heat
Exists in 2 forms:
Kinetic energy – energy of motion
Potential energy – energy at rest or energy of position

6. Energy
A roller coaster at the top of a hill has a great amount of potential energy.
As the rollercoaster begins to speed down the hill, the potential energy is turned into kinetic energy

7. Energy The SI unit for energy is the joule (J) 1 J = 1 Kgm2 / s2
Another unit of energy that you may be more familiar with is the calorie
calorie – amount of energy required to raise 1 g of water 1°C
1 cal = 4.18 J
1000 cal = 1Cal or kcal

8. Energy Conversions Convert 15,500 joules into Calories
15500 J x 1 cal x 1 Cal =
4.18 J cal
3.71 Cal

9. Calorimetry
A calorimeter is an insulated device used for measuring the amount of heat absorbed or released during a chemical or physical process.

10. Heat Calculations q = mcT q = heat (j) m = mass (g)
c = specific heat (j/gC)
T = change in temperature = Tf-Ti (C)

11. Heat Calculations
Specific Heat (c) – the amount of heat required to raise 1 gram of a substance by 1C
Specific heat is an intensive property
Every substance has its own specific heat

12. Heat Calculations
Specific Heat (c) – the amount of heat required to raise the temperature of 1 g of water 1 °C
The units for specific heat are J/g°C
The specific heat of water (in a liquid form is 4.18 J/g°C)
All substances have a particular specific heat

13. Specific Heat
A 10.0 g sample of iron changes temperature from 25.0C to 50.4 C while releasing 114 joules of heat. Calculate the specific heat of iron.

14. Example
q = mcT
114 = (10.0)c( )
c = j/g C

15. Another example
If the temperature of 34.4 g of ethanol increases from 25.0 C to 78.8 C how much heat will be absorbed if the specific heat of the ethanol is 2.44 j/g C

16. Another example
q = mcT
q = (34.4)(2.44)( )
q = 4.53 x 10 3 j

17. Yet another example
4.50 g of a gold nugget absorbs 276 J of heat. What is the final temperature of the gold if the initial temperature was 25.0 C & the specific heat of the gold is 0.129j/g C

18. Yet another example q = mcT 276 = (4.50)(0.129)T T = 475C
T = Tf-Ti
475 = Tf-25
Tf = 500 C

19. Endothermic & Exothermic Reactions
Endothermic reactions – chemical reaction that requires energy to break existing bonds
Heat goes into the reaction from the surroundings
Exothermic reactions – chemical reaction in which energy is released
Heat goes out of the reaction into the surroundings

20. Endothermic & Exothermic Reactions
Endothermic reactions – since heat goes from the surroundings into your system, it will feel colder
Temperature of endothermic reactions goes down
The sign for the heat change will be +
Exothermic reactions – since heat goes from the system to the surroundings, it will feel hotter
Temperature of exothermic reactions goes up
The sign for the heat change will be -

21. Endothermic & Exothermic Reactions

22. Endothermic & Exothermic Reactions

23. Endothermic & Exothermic Reactions
Are the following reactions endothermic or exothermic?
CO + 3H2  CH4 + H2O H= -206kJ
Exothermic (H is negative)
I add magnesium metal to some hydrochloric acid. The temperature goes from 23C to 27 C
Exothermic – temperature goes up
I mix together some vinegar & baking soda. The temperature goes from 28C to 23C
Endothermic – temperature goes down

24. Calorimetry
When using calorimetry, you are usually trying to determine the identity of an unknown metal by finding its specific heat
The heat lost from the metal will be gained by the water
q metal = - q water

25. Calorimetry q metal = - q water
(m metal)(c metal)(T metal) = - (m water)(c water)(T water)

26. Calorimetry Examples
A 58.0 g sample of a metal at °C is placed in a calorimeter containing 60.0 g of water at 18.0 °C. The temperature of that water increases to 22.0 °C. Calculate the specific heat of the metal.

27. Calorimetry Examples q metal = - q water
(m metal)(c metal)(T metal) = - (m water)(c water)(T water)
(58.0)(c metal)( ) = - (60.0)(4.18)( )
-4524 (c metal) = -1004
(c metal) = J/g°C

28. Calorimetry Examples
A piece of metal with a mass of 4.68 g at 135°C is placed in a calorimeter with 25.0 g of water at 20.0 °C. The temperature rises to 35.0 °C. What is the specific heat of the metal?

29. Calorimetry Examples q metal = - q water
(m metal)(c metal)(T metal) = - (m water)(c water)(T water)
(4.68)(c metal)( ) = - (25.0)(4.18)( )
-468 (c metal) =
(c metal) = 3.35 J/g°C

30. Specific Heat & Phase Changes

31. Changes in State
Increasing or decreasing the amount of kinetic energy will cause changes in the state of matter
Changes of State

32. Phase changes
Transitions between solid, liquid, and gaseous phases typically involve large amounts of energy compared to the specific heat. If heat were added at a constant rate to a mass of ice to take it through its phase changes to liquid water and then to steam, the energies required to accomplish the phase changes (called the latent heat of fusion and latent heat of vaporization ) would lead to plateaus in the temperature vs time graph. The graph below presumes that the pressure is one standard atmosphere.