Thermochemistry

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Chapter 16

• Thermochemistry

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• Thermochemistry the study of the transfers of
  energy as heat that accompany chemical reactions
  and physical changes.
• Temperature measure of the average kinetic
  energy of the particles of a substance.
• Heat energy transfer due to a difference in
  temperature.

Units Temperature Fahrenheit, Celsius,
Kelvin Heat Joule, Calorie
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Calorimetry

• Specific Heat amount of energy needed to raise
  the temperature of one gram of a substance one
  degree Celsius (or Kelvin).
• Units J/gC or J/gK

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Calorimeters
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How much heat is needed to change the temperature
of a 50 g piece of aluminum from 10 C to 80 C? q
mcDT q (50g)(.897J/gK)(80-10) q 3139.5
J What is the final temperature of a 50 g piece
of aluminum loses 8000 J of heat, if its initial
temperature was 90 C? q mc(Tf Ti) -8000 J
(50g)(.897J/gC)(Tf 90) -8000J /
(50g)(.897J/gC) Tf 90 -178.37C Tf 90 Tf
-178.37 90 -88.37C
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What is the final temperature of a coffee cup
system if a 50 g aluminum spoon at 15 C is placed
into a 250 mL(g) cup of coffee at 60 C? Heat Lost
(-)Heat Gained mcDTlost -mcDTgained (250g)(4.1
84J/gC)(Tf 60)-(50g)(.897J/gC)(Tf
15) 1046(Tf 60)-44.85(Tf 15) 1046Tf
62760-44.85Tf 672.75 1046Tf 44.85Tf 62760
672.75 1090.85Tf 63432.75 Tf
63432.75/1090.85 Tf 58.15 C
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Enthalpy of Reaction quantity of heat energy
transferred in a reaction (heat of reaction)

• DH Hproducts Hreactants

Thermochemical Equations 2H2(g) O2(g) ?
2H2O(g) 483.6 KJ Amount of energy is
proportional to the number of moles 4H2(g)
2O2(g) ? 4H2O(g) 967.2 KJ H2(g) 1/2O2(g) ?
H2O(g) 241.8 KJ Reverse reactions energy
changes sides 2H2O(g) 483.6 KJ ? 2H2(g)
O2(g)
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2H2(g) O2(g) ? 2H2O(g) 483.6 KJ is more
commonly seen as 2H2(g) O2(g) ? 2H2O(g) DH
-483.6 KJ 2H2O(g) 483.6 KJ ? 2H2(g)
O2(g) is more commonly seen as 2H2O(g) ? 2H2(g)
O2(g) DH 483.6 KJ -DH exothermic heat
released DH endothermic heat absorbed
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• Molar Enthalpy of Formation DHf enthalpy
  change that occurs when one mole of a compound is
  formed from its elements in their standard state
  _at_ 25 C and 1 atm.
• Elements in their standard state DHf 0
• DHf O2 or Ag or H2 0
• DHf CO2 -393.5 KJ/mol (-) means more
  stable than free elements
• DHf C2H2 226.7 KJ/mol () means reacts easily
  or is less stable than the free elements

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• Molar Enthalpy of Combustion DHc enthalpy
  change that occurs when one mole of a substance
  completely combusts.
• Combustion complete synthesis reaction with
  oxygen.
• H2(g) ½ O2(g) ? H2O(g) DHc -241.8 KJ/mol
• Appendix Table A-5

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• Hesss Law the overall enthalpy change in a
  reaction is the sum of the enthalpy changes of
  all the individual steps that make up the
  reaction.

• Consider the combustion reaction of methane to
  form CO2 and liquid H2O
• CH4(g) 2O2(g) -gt CO2(g) 2H2O(l)
• This reaction can be thought of as occurring in
  two steps
• In the first step methane is combusted to produce
  water vapor
• CH4(g) 2O2(g) -gt CO2(g) 2H2O(g)
• In the second step water vapor condenses from the
  gas phase to the liquid phase
• 2H2O(g) -gt 2H2O(l)

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Each of these reactions is associated with a
specific enthalpy change CH4(g) 2O2(g) -gt
CO2(g) 2H2O(g) DH -802 kJ 2H2O(g) -gt 2H2O(l)
DH -88 kJ
Combining these equations yields the
following CH4(g)2O2(g)2H2O(g) -gt
CO2(g)2H2O(g)2H2O(l) DH (-802) kJ (-88)
kJ -890 kJ
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Enthalpy of Reaction from Combustion Enthalpies

• C(s) 2H2(g) ? CH4(g)
• C(s) O2(g) ? CO2(g)
• H2(g) ½ O2(g) ? H2O(l)
• CH4(g) 2O2(g) ? CO2(g) 2H2O(l)
• C(s) O2(g) ? CO2(g)
• 2H2(g) O2(g) ? 2H2O(l)
• CO2(g) 2H2O(l) ? CH4(g) 2O2(g)

• DHc ?
• DHc -393.5 kJ
• DHc -285.8 kJ
• DHc -890.8 kJ
• DHc -393.5 kJ
• DHc (2)-285.8 kJ
• DHc 890.8 kJ

C(s) 2H2(g) ? CH4(g)
DHc -74.3kJ
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Enthalpy of Reaction from Formation Enthalpies
NO(g) ½ O2(g) ? NO2(g) DHrxn S DHf NO2(g) S
DHf NO DHf ½ O2(g) DHrxn 33.2kJ
90.29kJ 0 DHrxn - 57.1kJ
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Driving Force of Reactions

• Two factors that determine spontaneity
• Energy change
• Randomness of particles
• Most reactions are exothermic
• Some endothermic if there is an increase in
  randomness

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Entropy

• Entropy S measure of the degree of randomness
  of the particles in a substance.
• DS difference between the entropy of the
  products and reactants
• DS Sproducts Sreactants
• DS increase in entropy
• -DS decrease in entropy

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Free Energy

• Gibbs Free Energy G combined enthalpy
  entropy function.
• Natural processes proceed in a direction that
  lowers the free energy of a system
• DG DH TDS
• If DG lt 0 reaction is spontaneous

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Homework

• Pages 552-553
• Numbers 4,7,8,9,11,14,116,23,25