Transcript Enthalpies of Formation enthalpy of formation, H , or heat of formation

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) --> C2H5OH(l)
DHfo = -277.69 kJ
The standard enthalpy of formation of the most stable form of
an element under standard conditions is ZERO.
O2 (g) --> O2 (g)
DH = 0
1/2 N2 (g) + 3/2 H2 (g) --> NH3 (g) DHof = -46.19 kJ/mol
Using Enthalpies of Formation to calculate
Standard Reaction Enthalpies
Combustion of propane (C3H8) gas to form CO2(g) and H2O(l)
C3H8 (g) + 5 O2 (g) --> 3CO2 (g) + 4H2O(l)
This equation can be written as the sum of the following
three equations
C3H8(g) --> 3C(s) + 4H2(g)
DH1 = - DHfo (C3H8(g) )
+ 3C(s) + 3O2(g) --> 3CO2(g)
DH2 = 3 x DHfo (CO2(g) )
+ 4H2(g) + 2O2(g) --> 4H2O(l)
DH3 = 4 x DHfo (H2O (l) )
C3H8 (g) + 5 O2 (g) --> 3CO2 (g) + 4H2O(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
In general,
DHorxn = S n DHfo (products) - S n DHfo (reactants)
n is the stoichiometric coefficients in the reaction
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)
C6H6 (l) + 15 O2 (g) --> 6CO2 (g) + 3H2O(l)
2
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
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) => DHorxn = - 41.8 kJ/g
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) --> 2 H
DHB = 436 kJ/mol
DHo = +436 kJ
Mean bond enthalpy: average molar enthalpy change
accompanying the dissociation of a given type of bond.
Estimate the enthalpy change of the reaction between
gaseous iodoethane and water vapor.
CH3CH2I(g) + H2O(g) --> 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
Fuels
During the complete combustion of fuels, carbon is
completely converted to CO2 and hydrogen to H2O.
C3H8 (g) + 5 O2 (g) --> 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.
Hubbert’s Peak, K. S. Deffeyes
US crude oil production
Global Energy Reserves (1988) (units of Q = 1021 J)
Fuel Type
Proven Reserves
Est. Reserves
Coal
25Q
118Q
Oil
5Q
9Q
Natural Gas
4Q
10Q
Total amount of commercially energy currently consumed by
humans ~ 0.5Q annually
“Non-renewable” sources of energy
Alternate Fuels
Natural Gas and Propane
C(s) + O2(g) --> CO2(g) DH = -393.5 kJ/mol
CH4(g) + 2 O2(g) --> CO2(g) + 2 H2O(l) DH = -890 kJ/mol
C3H8(g) + 5 O2(g) --> 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
Name
C(s)
CH4(g)
C3H8(g)
Heat released per gram
34 kJ
55.6 kJ
50.3 kJ
Name
C(s)
CH4(g)
C3H8(g)
Heat released per mole of CO2(g) released
393.5 kJ
890 kJ
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
Methanol & Ethanol
Alcohols have the advantage over natural gas in that they
are liquids at atmospheric pressure and temperature.
Compound
CH3OH(l)
C2H5OH (l)
CH4(g)
C(s)
DHcombustion (kJ/g)
-22.7
-29.7
-55.6
-34
Hydrogen
H2(g) + 1/2O2(g) -------> H2O(l)
spark
DH = -286 kJ/mol
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
Compound
DHoc
Hydrogen (H2(g))
Methane (CH4(g))
Octane (C8H18(l))
Methanol (CH3OH(l))
kJ/mol
-286
-890
-5471
-726
Specific
Enthalpy
kJ/g
-142
-55
-48
-23
Enthalpy
density
kJ/L
-13
-40
-3.8 x 104
-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.
Spontaneous Change
A spontaneous change is one that occurs without external
intervention and has definite direction.
Spontaneous for
T > 0oC
Spontaneous for
T < 0oC
A spontaneous process need not be fast
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.
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.
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.
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.
Ice melts spontaneously at T>0oC 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.