Cem1000w - Thermodynamics notes

THERMODYNAMICS

Thermodynamics: the study of heat and its transformations
 System: the chemical reaction under investigation
 Surroundings: rest of the universe
 Heat: the energy which flows into or out of a system because of a difference in
temperature between the system and surroundings (transfers until thermal
equilibrium is established)
First Law of Thermodynamics

“Energy is neither created nor destroyed, it is conserved.”
“The total energy of the universe is constant.”
∆ E universe=∆ E system+∆ E surroundings
Internal energy

 ‘The change in the internal energy of a closed system ∆ E is equal to the sum
of the heat transferred into or out of the system and work done by or on the
system.’
 ∆ E =q+ w
 Heat transferred into system: q>0 (endo)
 Heat transferred out system: q<0 (exo)
 Work done on system by surroundings: w>0
 Work done by system on surroundings: w<0
Second Law of Thermodynamics
 Spontaneous process: A process that proceeds on its own without any outside
assistance (occurs in a definite direction)
 Processes that are spontaneous in one direction are spontaneous in the other.
Predicting Spontaneous Change

 The sign of ∆ H rxn. cannot be used to predict spontaneity as spontaneous
reactions may be exothermic OR endothermic.
 Considerations based on energy/enthalpy are inadequate for predicting spontaneity.
 The thermodynamic function which is an indicator of spontaneous change is the
state function called ENTROPY
 Entropy, S, is a measure of the extent of disorder in a system (units J.K-1)
 A system’s entropy is reflected in the specific distribution of the molecules among
the discrete energy levels available to them at a given temperature.
Approaches for quantifying changes in entropy of a system:

1. Statistical microstate approach
 Boltzmann equation: S = k ln(W)
 Calculate W in initial and final states to get ∆ S sys

, 2. Using standard molar entropies
 ∆ S ° = ∑ m S ° ( products ) −∑ n S ° ( reactants)
3. Heat transfer approach at constant temperature
 ∆ S =q/ T
 Determine q in initial and final states to get ∆ S sys
Microstates and Dispersal Energy
 A system of molecules has different allowed energy states
 Each quantized energy state for a system is called a microstate
 At any instant, the total energy of the system is dispersed throughout one microstate.
 Under a given set of conditions, each microstate has the SAME TOTAL energy as any
other.
 Each microstate is therefore, equally likely.
 The larger the number (W) of possible microstates, the larger the number of ways in
which a system can disperse its energy.
Spontaneous expansion of gas

 Increasing volume
 Number of translational energy levels increases
 Greater dispersal of energy among available energy levels
 Increased number of ‘microstates’
 Higher entropy

Opening the stopcock increases the number of possible energy levels, which are closer
together on average.
More distributions of particles are possible, therefore entropy increases
W final
= 2NA
W initial
∆ S sys=S final−S initial=klnW final−kln W initial

Where: W = N!/(n1!n2!...)

k=1.38x10-23J/K
Entropy

 A system with relatively few equivalent ways of arranging the components (small W)
has relatively ‘less disorder’ and low entropy
 A system with many equivalent ways of arranging the components (larger W) has
relatively ‘more disorder’ and high entropy
 Criterion for spontaneous change: Second Law of Thermodynamics
“The total entropy of the universe increases in any spontaneous process.”
∆ S universe= (∆ S system+∆ S surroundings )> 0
 The only direction in which a spontaneous change can occur is that for which the
total entropy of the universe increases.