1. Chapter 6: Thermochemistry
Thermodynamics: the Science of the Relationships Between Heat and Other Forms of Energy.
Thermochemistry: One Area of Thermodynamics That Involves the Quantity of Heat Absorbed or Evolved.
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2. Energy The capacity to do work or to produce heat.
Conservation of energy – energy can be converted from one form to another, but cannot be created nor destroyed.
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3. Forms of Energy Kinetic Energy Potential Energy
= ½ mv2
Energy of motion
Potential Energy
= mgh (physics)
Energy of position
Conversion between KE and PE
Ex) Tossing a stone into the air (LEP #1)
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5. KE and PE in Chemistry KE PE
Found in the various motions and vibrations of the molecules. PE
Found in the electrostatic attractions between atoms and ions (chemical bonds and intermolecular forces).
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6. Units of Energy SI unit is the joule (J)
= 1 kg m2 / s2
Relatively small unit – will usually use kilojoules (kJ)
1 kJ = 1000 J
Commonly used non-SI unit is the calorie (cal)
= amount of heat required to raise 1g of water by 1oC.
1 cal = 4.184J (exact)
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7. Energy Changes
Surroundings
System
System = the substances undergoing a chemical reaction or a physical change.
Surroundings = everything else.
Universe = System + Surroundings
What we study is the exchange of energy between the system and the surroundings
Surroundings
System
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8. Examining Energy Changes
Example: Adding NaOH(aq) to HCl(aq) in a styrofoam cup.
System = NaOH, HCl, H2O, and NaCl.
Surroundings = cup and everything else.
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9. Transferring Energy
Energy may be transferred in one of two forms or a combination of both.
Work = Force times distance, where a force is any push or pull on an object.
Heat = energy transferred from a hotter object to a colder one.
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10. Transferring Energy
All reactions in the lab thus far have been ones that transfer energy in the form of heat.
How can we have a reaction that does work?
A car engine does both work and heat.
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11. Sign Conventions
Work (w) and heat (q) must be defined from the direction of the system (and surroundings).
W = positive if work is done by the surroundings on the system and negative if work is done by the system on the surroundings.
q = positive if heat is transferred from the surroundings to the system and negative if heat is transferred from the system to the surroundings.
See Table 5.1 (p. 171) for summary.
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12. Internal Energy
The internal energy (E) of a system is defined as the sum of both the KE and PE of all the particles of a system.
Calculating E for any system is a rather daunting task.
When a system undergoes a physical or chemical change, the change in internal energy can be easily found.
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13. Internal Energy DE = q + w
Combustion of gasoline molecules in a car engine do work on the surroundings and transfer heat to the surroundings.
Thus, the internal energy of the system decreases.
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14. Exo and Endo
______________ reactions are reactions that transfer heat from the system to the surroundings.
______________ reactions are reactions that transfer heat from the surroundings to the system.
What signs (+/-) would be assigned to these definitions?
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15. State Functions
Any property that depends only on its present state and not the past states is called a state function (property).
A state function is independent of the path taken.
Ex) Change in altitude is a state function.
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16. State Functions
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17. First Law of Thermodynamics
The energy of the universe is constant
In a chemical process, energy can neither be created nor destroyed – that is energy is conserved.
Energy lost by the system is gained by that of the surroundings (and vice versa).
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18. Enthalpy Enthalpy (H) is a state function.
Enthalpy is a measurement of the flow of heat in a reaction.
Enthalpy is an extensive property.
DH = DE + PDV.
If no work is done, then DH = DE.
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19. Thermochemical Equation
A reaction with phase labels can now include the quantity of heat as a DH.
Ex) N2(g) + 3H2(g)  2NH3(g) ; DH = -92kJ
LEP #2
2 NH3(g)  N2(g) + 3H2(g) ?
2N2(g) + 6H2(g)  4NH3(g) ?
0.50 moles of NH3(g)
0.34g NH3(g)
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20. Calorimetry
A simple experiment for determining the change in enthalpy.
Measure the temperature change in an insulated cup.
Specific heat.
q = msDT
Heat capacity
q = CDT
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21. Calorimetry Calculating the specific heat of a metal (LEP #3)
Calculating the change in enthalpy using a specific heat (LEP #4 and #5)
Calculating the change in enthalpy using a heat capacity (LEP #6) by a bomb calorimeter
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22. Hess’s Law
If a reaction can be written as a series of steps that sum to the overall reaction, then the change in enthalpy for each step will sum to the overall change in enthalpy for the overall reaction.
Problem: can not directly measure a change in enthalpy because of issues with a reaction.
Ex) 2 C(graphite) + O2(g)  2 CO(g)
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23. Hess’s Law Can measure these:
C(graphite) + O2(g)  CO2(g) ; DH = kJ
2 CO(g) + O2(g)  2 CO2(g); DH = kJ
By manipulating these two equations, we can get them to add to the desired equation…
LEP #7
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24. Enthalpies of Formation
A set of standard values for various substances can be found in Appendix C.
These represent the formation of the substance from the elements in their standard states at 25oC and 1 atmosphere of pressure.
Ex) Na(s) + ½ Cl2(g)  NaCl(s) ; DHfo = kJ
Ex) C(gr) H2(g)  CH4(g) ; DHfo = kJ
Any element in its standard state will have a DHfo of zero.
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25. Enthalpies of Formation
To find a DH for a reaction, simply sum all the products and the reactants. Then subtract the reactants from the products.
DH = S(DHfo[products]) – S(DHfo[reactants])
Biggest source of error – sign mistakes!
Two problems –
Find a DH for an overall reaction – LEP#8.
Find a DHf0 for a compound in the reaction – LEP#9
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26. Foods and Fuels Foods provide energy for our bodies.
A Fuel is any substance that is burned to produce energy.
Values for each are compared by calculating the amount of energy per gram of substance.
This is also known as the Fuel Value.
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27. Foods Foods are made up of Carbohydrates, Proteins, and Fats.
Each of these has a caloric value.
Food labels list the breakdown of each of these three components in grams.
A Nutritional Calorie (note the capital “C”) is actually a kilocalorie.

28. Food Labels

29. Uniform Labeling
In 1990, the NLEA was passed to require that food labels contain certain information.
% Daily Value – reflects percents based on a 2,000 Calorie diet.
Good resource for finding caloric contents of foods including fast foods can be found at:
Serving sizes are set by the FDA from survey data.

30. Carbohydrates
These are a variety of compounds that include sugars and starches.
Glucose is the carbohydrate that is oxidized in our bodies for energy.
The net reaction is:
C6H12O6(s) + 6O2(g)  6CO2(g) + 6H2O(l)
and their DH’s from Appendix C are:
kJ, 0kJ, kJ, and kJ, respectively

31. Carbohydrates Thus, the DH for this reaction is calculated as:
DH = [(6)(-393.5kJ) + (6)(-285.8kJ)] – [(1)( kJ)]
DH = kJ/mol
This value, in nutritional terms, is converted to an amount of calories per gram of glucose.

32. Carbohydrates In nutrition, the negative sign is ignored.
Thus, our value is about 3700 cal or 3.7kcal.
But, this is just one of many examples of carbohydrates.
Average of all of carbohydrates is approximately 4 kcal/g (or 4 Cal/g).

33. Fats
A fat is a triglyceride made up of glycerol and three fatty acids.
Fatty acids are long chains of carbon atoms with the carboxyl group (COOH) at the end.
Animal fats contains mainly Stearic Acid, which has the formula C18H36O2.
Thus, Animal fat is called Tristearin, with a formula of C57H110O6.

34. Fats The balanced reaction is: The reaction generates a DH = -75,520kJ
2C57H110O6(s) + 163O2(g)  114CO2(g) + 110H2O(l)
The reaction generates a DH = -75,520kJ
The average caloric value for fats is about 9kcal/g.

35. Proteins
These are metabolized differently than a combustion reaction in our bodies.
Proteins are constructed from the 20 amino acids into very long chains.
The average caloric value for proteins is 4 kcal/g.

36. Calculating the Calories
The food label earlier contained 31g of carbohydrate, 12g of fat, and 5g of protein per serving.
Thus, the total Calories are:
31g x 4kcal/g
+ 12g x 9kcal/g
+ 5g x 4kcal/g
= 252 kcal or 252 Cal

37. Fuels
We use a variety of fuels to provide work or transfer heat energy.
Coal – burned to produce electricity
C(gr) + O2(g) CO2(g) ; DH = kJ
Fuel Value = kJ x 1mole C / 12.0g
= 32.8 kJ/g
Natural Gas – CH4(g) – used to heat our homes and water
CH4(g) + 2 O2(g) CO2(g) + 2 H2O(l)
DH = -802 kJ

38. Fuels Fuel Value = 802 kJ x 1mole CH4 / 16.0g = = 50.1 kJ/mol
Gasoline – a blend of hydrocarbons whose average formula is C8H18
2 C8H O2(g) CO2(g) + 18 H2O(l)
DH = -10,148 kJ
Fuel Value = 10,148 kJ/2mol x 1mol/114g
= 44.5 kJ/mol

39. Fuel Value So far… Is there any fuel that is better?
Coal < Gasoline < Natural Gas
Is there any fuel that is better?

40. Hydrogen! 2 H2(g) + O2(g) 2 H2O(g) DH = -484 kJ
Fuel Value = 484 / 2mol x 1mol/2.0g
= 121 kJ/g

41. Fuel Cells
A fuel cell combines hydrogen and oxygen to make water and electricity.
Pros and cons.

42. Integrative Exercises
Combine two (or more) parts of problems.
Ex) Combine calorimeter problem with finding an enthalpy of formation (LEP #10).
Ex) Using a q = msDT with a thermochemical equation (LEP #11).
Ex) Determine the final temperature of a solution (LEP #12).
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