Enthalpies of Formation

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• Enthalpies of Formation
• The enthalpy of formation, DHf, or heat of
  formation, is defined as the change in enthalpy
  when one mole of a compound is formed from its
  stable elements.
• The standard enthalpy of formation (DHfo) of a
  compound is defined as the enthalpy change for
  the reaction that forms 1 mole of compound from
  its elements, with all substances in their
  standard states.

2C(s) 1/2 O2(g) 3 H2 (g) --gt C2H5OH(l) DHfo
-277.69 kJ
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• The standard enthalpy of formation of the most
  stable form of an element under standard
  conditions is ZERO.
• O2 (g) --gt O2 (g) DH 0

1/2 N2 (g) 3/2 H2 (g) --gt NH3 (g) DHof -46.19
kJ/mol
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Using Enthalpies of Formation to calculate
Standard Reaction Enthalpies
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• Combustion of propane (C3H8) gas to form CO2(g)
  and H2O(l)

C3H8 (g) 5 O2 (g) --gt 3CO2 (g) 4H2O(l)
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This equation can be written as the sum of the
following three equations
C3H8(g) --gt 3C(s) 4H2(g) DH1 - DHfo
(C3H8(g) )
3C(s) 3O2(g) --gt 3CO2(g) DH2 3 x DHfo
(CO2(g) )
4H2(g) 2O2(g) --gt 4H2O(l) DH3 4 x DHfo
(H2O (l) )
DHorxn DH1 DH2 DH3
Looking up the standard heats of formation for
each equation DHorxn -(-103.85) 3(-393.5)
4(-285.8)) -2220 kJ
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In general, DHorxn S n DHfo (products) - S n
DHfo (reactants) n is the stoichiometric
coefficients in the reaction
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Calculate the standard enthalpy change for the
combustion of 1 mole of benzene (C6H6 (l)) to
CO2(g) and H2O(l). Compare the quantity of heat
produced by the combustion of 1.00 g of propane
(C3H8(g)) to that produced by 1.00 g of C6H6 (l)
First write a balanced equation for the
combustion of 1 mole of C6H6 (l)
DHorxn 6 DHfo(CO2) 3DHfo(H2O) -
1DHfo(C6H6)
(15/2)DHfo(O2)
6(-393.5 kJ) 3(285.8 kJ) - 49.0 kJ - 7.5(0
kJ) -3267 kJ
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For the combustion of 1 mole of propane DHorxn
-2220 kJ Hence for 1.00g propane, which
corresponds to 0.0227 mol propane, DHorxn
0.0227mol x -2220 kJ/mol - 50.3 kJ/g For C6H6
(l) gt DHorxn - 41.8 kJ/g
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Bond Enthalpies
Strength of a chemical bond is measured by the
bond enthalpy, DHB Bond enthalpies are positive,
because heat must be supplied to break a
bond. Bond breaking is endothermic Bond formation
is exothermic. H2(g) --gt 2 H DHo 436
kJ DHB 436 kJ/mol
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Mean bond enthalpy average molar enthalpy change
accompanying the dissociation of a given type of
bond.
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Estimate the enthalpy change of the reaction
between gaseous iodoethane and water
vapor. CH3CH2I(g) H2O(g) --gt CH3CH2OH(g)
HI(g) Reactant break a C-I bond and an O-H
bond DHo 238 kJ 463 kJ 701 kJ Product to
form a C-O bond and an H-I bond DHo -360 kJ
-299 kJ -659 kJ Overall enthalpy change 701
kJ - 659 kJ 42 kJ
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Fuels During the complete combustion of fuels,
carbon is completely converted to CO2 and
hydrogen to H2O. C3H8 (g) 5 O2 (g) --gt 3CO2
(g) 4H2O(l) Standard heats of formation of
CO2(g) and H2O(l) DHfo (CO2(g)) -393.5
kJ/mol DHfo(H2O(l)) -286 kJ/mol The greater
the percentage of carbon and hydrogen in a fuel,
the higher its fuel value.
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Hubberts Peak, K. S. Deffeyes
US crude oil production
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Global Energy Reserves (1988) (units of Q 1021
J) Fuel Type Proven Reserves Est.
Reserves Coal 25Q 118Q Oil 5Q 9Q Natur
al Gas 4Q 10Q Total amount of commercially
energy currently consumed by humans 0.5Q
annually Non-renewable sources of energy
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Alternate Fuels

• Natural Gas and Propane
• C(s) O2(g) --gt CO2(g) DH -393.5 kJ/mol
• CH4(g) 2 O2(g) --gt CO2(g) 2 H2O(l) DH -890
  kJ/mol
• C3H8(g) 5 O2(g) --gt 3CO2 4 H2O DH -2213
  kJ/mol

Natural gas, primarily methane with small amounts
of ethane and propane used for cooking and
heating. Highly compressed natural gas (CNG) -
commercial vehicles. Liquid petroleum gas (LPG)
- propane - also used as a fuel for vehicles
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Name Heat released per gram C(s) 34 kJ CH4(g)
55.6 kJ C3H8(g) 50.3 kJ
Name Heat released per mole of CO2(g)
released C(s) 393.5 kJ CH4(g) 890
kJ C3H8(g) 738 kJ
CH4(g) and C3H8(g) release more energy per gram
and can be considered to be cleaner
fuels. Disadvantages leakage of CH4 from pipes,
storage and transportation, need to be compressed
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• Methanol Ethanol
• Alcohols have the advantage over natural gas in
  that they are liquids at atmospheric pressure and
  temperature.
• Compound DHcombustion (kJ/g)
• CH3OH(l) -22.7
• C2H5OH (l) -29.7
• CH4(g) -55.6
• C(s) -34

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Hydrogen H2(g) 1/2O2(g) -------gt H2O(l) DH
-286 kJ/mol
spark
H2/O2 Fuel cells Electrical energy is produced
during the redox reaction
Advantages of using H2 as a fuel energy
released per gram low polluting Disadvantage
gas at room temperature
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Compound DHoc Specific
Enthalpy Enthalpy density
kJ/mol kJ/g kJ/L Hydrogen
(H2(g)) -286 -142 -13 Methane
(CH4(g)) -890 -55 -40 Octane
(C8H18(l)) -5471 -48 -3.8 x 104 Methanol
(CH3OH(l)) -726 -23 -1.8 x 104
Methane (CH4), Ethanol (C2H5OH), hydrogen (H2)
are renewable fuels. CH4 bacterial digestion
of waste H2 electrolysis of ocean water C2H5OH
biological fermentation of starches (e.g. in
corn) Combustion of CH4 and C2H5OH produce CO2,
but they produce less CO2 per gram than gasoline.
And they are renewable.
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Spontaneous Change

• A spontaneous change is one that occurs without
  external intervention and has definite direction.

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A spontaneous process need not be fast
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The change in enthalpy during a reaction is an
important factor in determining whether a
reaction is favored in the forward or reverse
direction. Are exothermic reaction more likely
to be spontaneous than an endothermic
reaction? Not necessarily. The endothermic
dissolution of ammonium nitrate, NH4NO3, occurs
spontaneously.
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• Entropy
• Both endothermic and exothermic reactions can be
  spontaneous
• Are there additional factors which determine
  spontaneity?
• Energy and matter tend to become more disordered.
• A measure of disorder is ENTROPY.

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When the valve is open, there are four possible
arrangements or STATES for both particles. Note
these arrangements are all equal in
energy. Opening the valve allows a higher degree
of disorder. The reverse process of the two gas
particles occupying only one flask is not
spontaneous.
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• As the number of particles increases in the
  system, the number of possible arrangements that
  the system can be in increases

Processes in which the disorder of the system
increases tend to occur spontaneously.
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• Ice melts spontaneously at Tgt0oC even though it
  is an endothermic process.
• The molecules of water that make up the ice
  crystal lattice are held rigidly in place.
• When the ice melts the water molecules are free
  to move around, and hence more disordered than in
  the solid lattice.
• Melting increases the disorder of the system.