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Chemical Reactions
Chemical Reaction - Observation
Reaction (1)
CH4 + 2O2 CO2 + 2H2O
Reaction (2)
CH4 + CO2 2CO + 2H2
When carrying out these reactions we found that

at 400K (123°C), the reaction (1) will proceed and reaction (2) will not

at 1000K(723°C), reactions (1) & (2) both proceed; but


rxn (1) can go complete (until CH4 or CO2 consumed completely)

rxn (2) won’t complete (with a feed CH4=CO2=1 & CO=H2=0, max. conv.=63% at 1000K)
The reaction (1) will give out heat, but the reaction (2) will require heat.
Why?
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Chemical Reactions
Chemical Reaction Thermodynamics

Each molecule contains certain types and quantity of chemical energy

There is always energy change In a chemical reaction because of

breaking / reformation of chemical bonds
 out-giving or in-taking heat

There are different energies associated with a substances & a reaction
(A systematic study of various forms of energy & their changes is called
Thermodynamics)
We will learn some of these energies

The meanings

How to get values / do simple calculate

How to use them as a tool to study chemical reactions
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Chemical Reactions
Chemical Reaction Thermodynamics

The heat of formation, H, (also called Enthalpy of Formation or Enthalpy)

H is an energy associated with heat
 H is specific for each substance and is dependent of temperature & pressure
e.g. at 1000K: H°CH4=-89, H°O2=0, H°CO2=-394, H°H2O=-241, H°CO=-111, H°H2=0
(kJ/mol)
(H values for various substances can be found in physical chemistry/Chem Eng handbooks)

In a reaction we are interested in the enthalpy change, DH, which is calculated using
DHT0  ( vi Hi0,T )prod  ( vi Hi0,T )reac
refers to standard pressure (1 atm.)
Temperature
For

Rxn(1) CH4 + 2O2 CO2 + 2H2O
DH°1000=-801 kJ/mol
Rxn (2) CH4 + CO2 2CO + 2H2
DH°1000=+260k J/mol
The meaning

When DH<0, a reaction releases heat  reaction is exothermic, as in rxn (1)

When DH>0, a reaction requires heat  reaction is endothermic, as in rxn (2)
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Chemical Reactions
Chemical Reaction Thermodynamics

The Gibbs Free Energy, G, (also called Free Energy)
G
is a thermodynamic function related to a reaction. It is a function of H, T & S (entropy)
G
is specific for each substance & is a function of H, T & S (entropy)
e.g. at 1000K: G°CH4=-+30, G°O2=0, G°CO2=-395, G°H2O=-192, G°CO=-200, G°H2=0
(kJ/mol)
(G values for various substances can be found in physical chemistry/Chem Eng handbooks)

The Gibbs Free energy change, DG, in a reaction can be calculated using
DGT0  ( viGi0,T )prod  ( viGi0,T )reac
For

Rxn(1) CH4 + 2O2 CO2 + 2H2O
DG°400 DG°1000=-801 kJ/mol
Rxn (2) CH4 + CO2 2CO + 2H2
DG°400=+145, DG°1000=-24 kJ/mol
Use of DG - Rxn(1) DG <0 at 400 & 1000K-spontaneous, Rxn(2) DG <0 at 400K, will not proceed
for a reaction at
constant T, P,
 DGT   0 reaction can proceed (but we don’t know how fast it will be!)

DGT   0 reaction at equilibrium (no further change possible-‘dead’ state)
DG   0 reaction will NOT proceed (or can proceed backward!)
 T
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Chemical Reactions
Chemical Reaction Thermodynamics
Example of DH°T calculation•
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
Coeff.
H°400(
H°1000
1
-77
-89
2
0
0
Equation to use
1
-393
-394
2
-242
-248
CH4 + CO2  2CO + 2H2
1
-77
-89
1
-393
-394
2
-110
-111
2
0
0
kJ/mol
kJ/mol
DHT0  ( vi Hi0,T )prod  ( vi Hi0,T )reac
DH°400 =[1x(-393)+2x(-242)]-[1x(-77)+2x(0)]= -800 kJ/mol
DH°1000 =[1x(-394)+2x(-248)]-[1x(-89)+2x(0)]= -801 kJ/mol
DH°400 =[2x(-110)+2x(0)]-[1x(-77)+1x(-393)]=+250 kJ/mol
Reaction (1)
Reaction (2)
DH°1000=[2x(-111)+2x(0)]-[1x(-89)+1x(-393)]= +260 kJ/mol
Note: The heat of formation of single element gases (O2, H2, N2 etc) is defined as zero.
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Chemical Reactions
Chemical Reaction Thermodynamics
Example of DH°T calculation•
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
Coeff.
1
2
1
2
CH4 + CO2  2CO + 2H2
1
1
2
2
G°400(
-42
0
-394
-224
-42
-394
-146
0
kJ/mol
G°1000 +19
0
-396
-193
+19
-396
-200
0
kJ/mol
Equation to use
DGT0  ( viGi0,T )prod  ( viGi0,T )reac
DG°400 =[1x(-394)+2x(-224)]-[1x(-42)+2x(0)]= -800 kJ/mol
DG°1000 =[1x(-396)+2x(-193)]-[1x(+19)+2x(0)]= -801 kJ/mol
DG°400 =[2x(-146)+2x(0)]-[1x(-42)+1x(-394)]=+144 kJ/mol
Reaction (1)
Reaction (2)
DG°1000=[2x(-200)+2x(0)]-[1x(19)+1x(-396)]= -23 kJ/mol
Note: The Gibbs Free Energy of single element gas (O2, H2, N2 etc) is defined as zero.
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Chemical Reactions
Chemical Reaction Thermodynamics

The values of DG°T and DH °T
Equations
DHT0  ( vi Hi0,T )prod  ( vi Hi0,T )reac
DGT0  ( viGi0,T )prod  ( viGi0,T )reac
In both cases G° and H° values for the reactants and products have to be those at the
reaction temperature, indicated by the subscript.

For common substances, G° and H° values are given as a function of T in handbooks - okay

For some less common substances, you may only find values at 298K, G°298 and H°298
How do you convert values of G°298 and H°298 to those of G°T and H°T ?
Here is the equations you can use to calculate the values of G°T and H°T from G°298 and H°298
H i0,T  H i0,298   DC p ,i dT   DH phasechange,i
T
298
Gi0,T  H i0,T - T Si0,T
T
0
where Si0,T  S 298
  DC p ,i
298
dT Q phase

T
Tj
in which, S°T is the entropy and Cp is the heat capacity at constant pressure
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Chemical Reactions
Chemical Reaction Thermodynamics

Summary

Will a reaction proceed in the direction specified?
 Check DG°T value of the reaction. The DG°T value of a reaction can be calculated by
DGT0  ( viGi0,T )prod  ( viGi0,T )reac
The G° values of reactants / products can be found in literature. Remember
for a reaction at
constant T, P,
DGT   0 reaction can proceed (but we don’t know how fast it will be!)

DGT   0 reaction at equilibrium (no further change possible-‘dead’ state)
DG   0 reaction will NOT proceed (or can proceed backward!)
 T

Is a reaction exothermic or endothermic?
 Check DH°T value of the reaction. The DH°T value of a reaction can be calculated by
DHT0  ( vi Hi0,T )prod  ( vi Hi0,T )reac
The H° values of reactants / products can be found in literature
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