Physical Chemistry
Lecture 24
Thermodynamics and Chemical
Reactions
Chemical reaction processes
Chemical reaction – a thermodynamic
process in which a system of atoms formed
as molecules changes state
1
O2 ( g ) → H 2O (l )
2
N 2 ( g ) + O2 ( g ) → 2 NO ( g )
H2 (g) +
4 S 2 ( g ) → S8 ( g )
General form of a chemical reaction
a A + b B +  → cC
+ dD + 
Stoichiometry numbers
Chemical reaction is determined, in part, by amount of
material undergoing the process
Stoichiometry number – number of moles of each
reactant or product in the reaction
Example:
H2 (g) +

1
O2 ( g ) → H 2O (l )
2
Stoichiometry numbers
 Hydrogen
 Oxygen
 Water
-1
-1/2
+1
Stoichiometry numbers have a sign


Positive for products
Negative for reactants
Determining thermodynamics
of reactions
Can measure directly
Can determine changes in any state function by
any route

Allows one to determine change without
measurement
Example: Aldol condensation
2 (H 3C )2 CO (l ) →
(CH 3 )2 C (OH )CH 2C (O )CH 3 (l )
Stoichiometry and
thermodynamics of reactions
Chemical reactions may be written in the
following manner
∑ν
k
Ak
= 0
k
The change in any thermodynamic quantity
(e.g. enthalpy) during reaction is
θ
∆H reaction
=
θ
ν
H
∑ k m,k
k
Determining thermodynamics
of reactions
Can determine changes in any state function by
any route
Example: Aldol condensation


Change reactants to elements in step 1
Change elements to products in step 2
2 (H 3C )2 CO (l )
1
5 C ( gr ) + 6 O2 ( g ) + 6 H 2 ( g )

→
2

→
(CH 3 )2 C (OH )CH 2C (O )CH 3 (l )
θ
∆H reaction
= ∆H1θ
+ ∆H 2θ
θ
∆Greaction
= ∆G1θ
+ ∆G2θ
Thermodynamics of real and
theoretical chemical reaction
Most real reactions involve change from a
mixture of reactants to a mixture of products
Changes of thermodynamic functions are
usually theoretically calculated, starting from
pure reactants and ending at pure products
Calculations therefore neglect the effects of
mixing

May, in certain circumstances, determine how the
reaction goes
Standard enthalpy of reaction at
298.15K and 1 bar
Reaction
Cl2 (g) → 2 Cl (g)
C (gr) → C (dia)
4 C (gr) + 2 H2 (g) → CH4 (g)
ΔHθ (kJ)
+243.36
+1.90
-74.81
C (gr) + ½ O2 (g) → CO (g)
-110.53
H2 (g) + ½ O2 (g) → H2O (g)
-241.82
H2 (g) + ½ O2 (g) → H2O (l)
-285.83
C (gr) + O2 (g) → CO2 (g)
-393.51
2 H2 (g) + O2 (g) → 2 H2O (l)
-571.66
Extent of reaction and
thermodynamic quantities
Systems have many states



Pure reactants
Pure products
Intermediate mixtures
States represented by extent of reaction, ξ
Thermodynamic change to any state
(starting from pure reactants) is less than
complete change
θ
θ
∆H (ξ ) = ξ ∆H reaction
Calorimetry
Experimental determination of heat given off
by a system undergoing a process
C
Determine the heat capacity of
a bath by measuring
temperature rise of a bath with
a standard process
Determine the temperature rise
when the same bath is heated
by a specific process whose
heat change must be found
Calculate the heat of the
process
heat of process ≈ C ∆T
≈
heat of standard process
temperature change
Calorimeters
Modern Parr
calorimeter
Traditional Parr bomb
calorimeter
Reaction catalogues –
reporting calorimetric results
Many different reactions studied over the
past 150 years
Listing all reactions ever studied is
cumbersome
Catalogues usually contain “distillation” of
information to a specific kind of reaction
(usually at a specific set of conditions)



Formation reactions
Combustion reactions
NIST evaluates all published thermodynamic
data and publishes a list of accepted values
(http://www.nist.gov/srd/thermo.htm)
Standard formation reaction
Reaction to create material in its standard
state from elements in their standard states
Elements → Compound
Usually reported on a molar basis
Examples
H 2 ( g ,1bar ) +
1
O2
2
→ H 2O (l )
3
O2 ( g ,1bar ) → O3 ( g ,1bar )
2
3 C ( gr ) + 4 H 2 ( g ,1bar ) → C3 H 8 ( g ,1bar )
C ( gr ) + 2 H 2 ( g ,1bar ) + S ( s, rh) → H 3CSH ( g ,1bar )
Standard enthalpies of formation
of alkanes at 298.15K
Gases
Compound
Liquids
∆fHmθ (kJ/mole)
Compound
∆fHmθ (kJ/mole)
CH4
-74.9
C2H6
-83.8
C3H8
-104.7
C3H8
-119.8
n-C4H10
-127.1
n-C4H10
-125.6
i-C4H10
-134.2
n-C5H12
-146.8
n-C5H12
-173.5
n-C6H14
-167.2
n-C6H14
-198.7
n-C7H16
-187.8
n-C7H16
-224.4
n-C8H18
-208.4
n-C8H18
-250.3
Standard combustion reaction
Reaction of compound with oxygen to
form stable oxides
Compound + Oxygen → Stable Oxides
Usually reported on a molar basis
Examples
1
O2 → H 2O (l )
2
C ( gr ) + O2 ( g ,1bar ) → CO2 ( g ,1bar )
H 2 ( g ,1bar ) +
H 3CSH ( gr ) + 3 O2 ( g ,1bar ) → 2 H 2O (l ) + CO2 ( g ,1bar ) + SO2 ( g ,1bar )
Summary
Chemical reactions are thermodynamic processes
State changes in chemical reaction characterized by
changes of thermodynamic quantities, ∆U, ∆H, ∆S,
∆A, etc.
Calorimetry measures temperature change for a
bath heated when a system undergoes a process

Measures heat by comparison to known processes
(calibration)
Catalogues of reactions between specific states


Standard formation reaction
Standard combustion reaction