1. Thermochemistry

2. The Nature of Energy Kinetic Energy and Potential Energy
Kinetic energy is the energy of motion:
Potential energy is the energy an object possesses by virtue of its position.
Potential energy can be converted into kinetic energy. Example: a bicyclist at the top of a hill.

3. The Nature of Energy Units of Energy
SI Unit for energy is the joule, J:
sometimes the calorie is used instead of the joule:
1 cal = J (exactly)
A nutritional Calorie:
1 Cal = 1000 cal = 1 kcal

4. Thermochemistry Terminology
System: part of the universe we are interested in.
Surrounding: the rest of the universe.
Boundary: between system & surrounding.
Exothermic: energy released by system to surrounding.
Endothermic: energy absorbed by system from surr.
• Work ( w ): product of force applied to an object over a distance.
Heat ( q ): transfer of energy between two objects

5. The First Law of Thermodynamics
Internal Energy
Internal Energy: total energy of a system.
Involves translational, rotational, vibrational motions.
Cannot measure absolute internal energy.
Change in internal energy,
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6. The First Law of Thermodynamics
Relating DE to Heat(q) and Work(w)
Energy cannot be created or destroyed.
Energy of (system + surroundings) is constant.
Any energy transferred from a system must be transferred to the surroundings (and vice versa).
From the first law of thermodynamics:

7. The First Law of Thermodynamics

8. First Law of Thermodynamics
Calculate the energy change for a system undergoing an exothermic process in which 15.4 kJ of heat flows and where 6.3 kJ of work is done on the system.
DE = q + w

10. The First Law of Thermodynamics
Exothermic and Endothermic Processes
Endothermic: absorbs heat from the surroundings.
An endothermic reaction feels cold.
Exothermic: transfers heat to the surroundings.
An exothermic reaction feels hot.

12. Endothermic Reaction Ba(OH)2•8H2O(s) + 2 NH4SCN(s) 
Ba(SCN)2(s) NH3(g) + 10 H2O(l)

13. The First Law of Thermodynamics
State Functions
State function: depends only on the initial and final states of system, not on how the internal energy is used.

14. Zn(s) + 2H+(aq)  Zn2+(aq) + H2(g)
Enthalpy
Chemical reactions can absorb or release heat.
However, they also have the ability to do work.
For example, when a gas is produced, then the gas produced can be used to push a piston, thus doing work.
Zn(s) + 2H+(aq)  Zn2+(aq) + H2(g)
The work performed by the above reaction is called pressure-volume work.
When the pressure is constant,

15. Enthalpy

16. Enthalpy
Enthalpy, H: Heat transferred between the system and surroundings carried out under constant pressure.
Enthalpy is a state function.
If the process occurs at constant pressure,

17. Enthalpy Since we know that We can write
When DH is positive, the system gains heat from the surroundings.
When DH is negative, the surroundings gain heat from the system.

18. Enthalpy => Heat of Reaction

19. Enthalpies of Reaction
For a reaction:
Enthalpy is an extensive property (magnitude DH is directly proportional to amount):
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) DH = -802 kJ
2CH4(g) + 4O2(g)  2CO2(g) + 4H2O(g) DH = kJ

20. Enthalpies of Reaction
When we reverse a reaction, we change the sign of DH:
CO2(g) + 2H2O(g)  CH4(g) + 2O2(g) DH = +802 kJ
Change in enthalpy depends on state:
H2O(g)  H2O(l) DH = -44 kJ

21. Calorimetry Heat Capacity and Specific Heat
Calorimetry = measurement of heat flow.
Calorimeter = apparatus that measures heat flow.
Heat capacity = the amount of energy required to raise the temperature of an object (by one degree).
Molar heat capacity = heat capacity of 1 mol of a substance.
Specific heat = specific heat capacity = heat capacity of 1 g of a substance.

22. Table 5.2: Specific Heats (S) of Some Substances at 298 K
S ( J g-1 K-1 )
N2(g)
1.04
Al(s)
0.902
Fe(s)
0.45
Hg(l)
0.14
H2O(l)
4.184
H2O(s)
2.06
CH4(g)
2.20
CO2(g)
0.84
Wood , Glass
1.76 , 0.84

23. If 24. 2 kJ is used to warm a piece of aluminum with a mass of 250
If 24.2 kJ is used to warm a piece of aluminum with a mass of 250. g, what is the final temperature of the aluminum if its initial temperature is 5.0oC?
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24. Calorimetry Constant Pressure Calorimetry
Atmospheric pressure is constant!

25. Calorimetry
Constant Pressure Calorimetry

26. Calorimetry
Aqueous Solution Assumptions:

27. Calorimetry Examples
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In an experiment similar to the procedure set out for Part (A) of the Calorimetry experiment, g of Mg(s) was combined with mL of 1.0 M HCl. The initial temperature was 25.0oC and the final temperature was 72.3oC. Calculate: (a) the heat involved in the reaction and (b) the enthalpy of reaction in terms of the number of moles of Mg(s) used. Ans: (a) –25.0 kJ (b) –406 kJ/mol
50.0 mL of 1.0 M HCl at 25.0oC were mixed with 50.0 mL of 1.0 M NaOH also at 25.0oC in a styrofoam cup calorimeter. After the mixing process, the thermometer reading was at 31.9oC. Calculate the energy involved in the reaction and the enthalpy per moles of hydrogen ions used. Ans: kJ , -58 kJ/mol [heat of neutralization for strong acid/base reactions]

28. In an experiment similar to the procedure set out for Part (A) of the Calorimetry experiment, g of Mg(s) was combined with mL of 1.0 M HCl. The initial temperature was 25.0oC and the final temperature was 72.3oC. Calculate: (a) the heat involved in the reaction and (b) the enthalpy of reaction in terms of the number of moles of Mg(s) used.

29. qrxn= - S x msoln x ∆T
∆H = -406 kJ/mol

30. 50. 0 mL of 1. 0 M HCl at 25. 0oC were mixed with 50. 0 mL of 1
50.0 mL of 1.0 M HCl at 25.0oC were mixed with 50.0 mL of 1.0 M NaOH also at 25.0oC in a styrofoam cup calorimeter. After the mixing process, the thermometer reading was at 31.9oC. Calculate the energy involved in the reaction and the enthalpy per moles of hydrogen ions used. Ans: kJ , -58 kJ/mol [heat of neutralization for strong acid/base reactions]

31. Hess’s Law
Hess’s law: if a reaction is carried out in a number of steps, H for the overall reaction is the sum of H for each individual step.
For example:
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) H = -802 kJ
2H2O(g)  2H2O(l) H= kJ
CH4(g) + 2O2(g)  CO2(g) + 2H2O(l) H = kJ

32. Another Example of Hess’s Law
Given:
C(s) + ½ O2(g)  CO(g) DH = kJ
CO2(g)  CO(g) + ½ O2(g) DH = kJ
Calculate DH for: C(s) + O2(g)  CO2(g)

33. Enthalpies of Formation
If 1 mol of compound is formed from its constituent elements, then the enthalpy change for the reaction is called the enthalpy of formation, Hof .
Standard conditions (standard state): Most stable form of the substance at 1 atm and 25 oC (298 K).
Standard enthalpy, Ho, is the enthalpy measured when everything is in its standard state.
Standard enthalpy of formation: 1 mol of compound is formed from substances in their standard states.

35. Enthalpies of Formation
If there is more than one state for a substance under standard conditions, the more stable one is used.
Standard enthalpy of formation of the most stable form of an element is zero.

36. Enthalpies of Formation
Substance
DHof (kJ/mol)
C(s, graphite)
O(g)
247.5
O2(g)
N2(g)

38. Bucky Ball drawn by HyperChem

39. Enthalpies of Formation C3H8(g) + 5O2(g)  3CO2(g) + 4H2O()
Using Enthalpies of Formation to Calculate Enthalpies of Reaction
For a reaction
Note: n & m are stoichiometric coefficients.
Calculate heat of reaction for the combustion of propane gas giving carbon dioxide and water.
C3H8(g) + 5O2(g)  3CO2(g) + 4H2O()

40. C3H8(g) + 5O2(g)  3CO2(g) + 4H2O()
∆Hof (kJ/mol):

41. C3H8(g) + 5O2(g)  3CO2(g) + 4H2O()
∆Hof (kJ/mol):
∆Horxn = -2,219.9 kJ/mol

42. Foods and Fuels
Foods
1 nutritional Calorie, 1 Cal = 1000 cal = 1 kcal.
Energy in our bodies comes from carbohydrates and fats (mostly).
Intestines: carbohydrates converted into glucose:
C6H12O6 + 6O2  6CO2 + 6H2O, DH = kJ
Fats break down as follows:
2C57H110O O2  114CO H2O, DH = -75,520 kJ
Fats contain more energy; are not water soluble, so are good for energy storage.

44. Foods and Fuels
Fuels
Fuel value = energy released when 1 g of substance is burned.
Most from petroleum and natural gas.
Remainder from coal, nuclear, and hydroelectric.
Fossil fuels are not renewable.
In 2000 the United States consumed 1.03  1017 kJ of fuel.
In 2005 the United States consumed 1.05  1017 kJ of energy.
Hydrogen has great potential as a fuel with a fuel value of 142 kJ/g. [ gasoline ≈ 35 kJ/g ]

45. Foods and Fuels [8.1%] [6.0%] [40.4%] [22.9%] (% for 2000)
[22.6%]

46. Energy in Fuels
The vast majority of the energy consumed in this country comes from fossil fuels.
© 2012 Pearson Education, Inc.

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Thermochemistry