Thermodynamics

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Thermodynamics

• The study of energy and its interconversions

First Law of Thermodynamics

• The energy of the universe is constant
• Energy can be converted from one form to another
  but can be neither created nor destroyed

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Energy Capacity to do work or to transfer heat
Kinetic Energy Energy of motion ½ mv2 (m
mass of object v velocity) Potential
Energy Energy due to position or composition
(Chemical Energy) Work Force acting over a
distance Heat Energy transferred from hotter
object to colder object
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Energy defined from systems point of view
?Esystem Esystem final Esystem initial
Thermodynamic quantities have a number and a sign
Surroundings
Surroundings
System
System
Energy
Energy
DE lt 0 negative
DE gt 0 positive
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Internal Energy (E)

• Sum of the kinetic energy and potential energy
  of all particles in the system
• Unit is Joule (J)
• 1 J kg m2 s-2
• Calorie (cal) English Unit
• 1 cal 4.184 J

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Internal Energy (E)

• Energy can be changed by a flow of work, heat or
  both
• DE q w
• DE change in energy
• q heat
• w work

?
sign into system - sign out of system
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Chemical Energy

• Exothermic Process
• Heat is evolved (q is negative)
• Energy flows out of system
• Endothermic Process
• Heat is absorbed (q is positive)
• Energy flows into system

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Enthalpy (H)
DH qp ? DE

• Heat of reaction is the same as the change in
  enthalpy
• At constant pressure, the change in enthalpy DH
  of the system is equal to the energy flow as heat
  (w ? 0)
• Enthalpy is a state function
• The value does not depend on the history of the
  sample (talk more about later)

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Chemical Reactions and DH
DH Hproducts Hreactants

• Exothermic Process
• Reactants have greater H than products
• Heat is evolved (DH is negative)
• Energy flows out of system (we feel heat)
• Endothermic Process
• Products have greater H than reactants
• Heat is absorbed (DH is positive)
• Energy flows into system (we feel cold)

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Enthalpy and moles

• Enthalpy is an extensive property
• ie. magnitude of DH is proportional of the amount
  of reactant consumed in the process
• Sign of enthalpy is defined for a particular
  reaction path

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Specific Heat CapacitySpecific Heat q / (m ?
?T)

• The energy required to raise the temperature of
  one gram of a substance by one degree celsius
• Units are J/(C g) also written J C-1g-1
• Molar heat capacity is in moles rather than grams
  (J C-1mol-1)

Specific Heat Examples H2O(l) 4.18 J/(C
g) Fe(s) 0.45 J/(C g)
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Which get hotter?

• Consider two 15 g blocks at 24C, one made out of
  iron and the other made out of aluminum. If you
  supply both blocks with 12 kJ of heat, which will
  be hotter? How much hotter?

Specific heat of Al(s) 0.90 J C-1g-1 Specific
heat of Fe (s) 0.45 J C-1g-1
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Calorimeter

• A device used to determine the heat associated
  with a chemical process
• Coffee-cup calorimeter shown is used for
  constant-pressure calorimetry
• Water absorbs heat from process
• We record temperature changes in water

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Calorimetry Calculations

• Exothermic Processes

Energy released by rxn energy absorbed by
solution solution gets hot

• Endothermic Processess

Energy absorbed by rxn energy released by
solution solution gets cold
E (?H) transferred by rxn E (?H) Absorbed by
Soln
opposite signs
specific heat of solution J /(C g)
mass of solution (g)
increase in temperature (C)
E (?H) Absorbed by Soln

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Energy Transferred per Mole

• Energy transferred depends on amount of limiting
  reagent (treat energy like a product with a
  coefficient of 1)
• Ethalpy (DH) expressed in J/mol or kJ/mol
• Sign of enthalpy depends on whether energy is
  released or absorbed

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Group Question
In a coffee-cup calorimeter, 1.21 g of CaCl2 is
mixed with 65.1 g of water at an initial
temperature of 24.00 C. After dissolution of
the salt, the final temperature of the
calorimeter contents is 26.00 C. Assuming the
solution has a heat capacity of 4.18 J/ (C g)
and assuming no heat loss to the calorimeter,
calculate the enthalpy change for the dissolution
of CaCl2 in kJ/mol. (Hint the solution mass is
equal to the mass of the water plus the mass of
the salt.)
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Hesss Law

• The change in enthalpy is not dependent on the
  pathway
• In going from reactant to products, the number
  of steps the reactant(s) take not effect DH

N2 (g) 2 O2(g) ? 2 NO2 (g) DH1 68 kJ
N2 (g) O2 (g)? 2 NO (g) DH2 180 kJ 2
NO (g) O2 (g) ? 2 NO2 (g) DH3 -112 kJ

N2 (g) 2 O2(g) ? 2 NO2 (g) DH2 DH3 68 kJ
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Using Hesss law

• Find reactions that add up to the one of
  interest, states of matter are important
• If a reaction is reversed, reverse the sign of DH
• If coefficients in a balanced equation are
  multiplied by an integer, DH is also multiplied

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Standard Enthalpy of Formation

• Change in enthalpy that accompanies the
  formation of one mole of a compound from its
  elements with all substances in their standard
  states
• Standard states reflect the elemental form
• ie. the standard state of carbon is C (s)
• ie. the standard state of oxygen is O2 (g)
• Standard enthalpies of formations are listed in
  Appendix C

½ N2 (g) O2 (g) ? NO2 (g) DHf 34 kJ/mol
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Using Standard Enthalpies

• Using Hesss law and standard enthalpies, the
  enthalpy of a reaction can be determined
  (estimated)

DHrxn ? np DHf (products) - ? nr DHf
(reactants)
n coefficient from balanced equation DHf
standard enthalpy of formation Note DHf 0
for all elements in standard states
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Using Hesss Law and Standard Enthalpies of
Formation