Thermodynamics
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Thermodynamics
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Kinetics
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Way to calculate if a reaction will occur
Way to determine the rate of reactions
Thermodynamic equilibrium rarely
attained:
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Biological processes – work against thermo
Kinetic inhibitions
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Thermodynamics very useful
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Good approximation of reactions
Tells direction a reaction should go
Basis for estimated rates
Farther from equilibrium, faster rate
Thermodynamic definitions
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System – part of universe selected for
study
Surroundings (Environment) – everything
outside the system
Universe – system plus surroundings
Boundary – separates system and
surroundings
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Real or imagined
Boundary conditions – solutions to Diff Eq.
Types of systems
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Open system
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Closed system
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Exchanges with surroundings
Mass, also heat and work
no exchange of matter between surrounding
and system, energy can be exchanged
Isolated system
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there is no interaction with surroundings, no
exchange of energy or matter
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Steady state system
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Flux in = flux out
There can be exchange, but no change in
total abundance
Parts of Systems
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Phase – physically and chemically
homogeneous region
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Example: saturated solution of NaCl
Species – chemical entity (ion, molecule,
solid phase, etc.)
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E.g. NaCl (solid) + H20 (liquid)
Also Na+, Cl-, OH-, H+, NaClo, others
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Components
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Minimum number of chemical entities required
to define compositions of all species
Many different possibilities
Na+, Cl-, H+, OH NaCl – H2O
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Thermodynamic Properties
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Extensive
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Depends on amount of material
E.g., moles, mass, energy, heat, entropy
Additive
Intensive
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Don’t depend on amount of material
Concentrations, density, T, heat capacity
Can’t be added
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State function
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a property of a system which has a specific
value for each state (e.g., condition)
E.g., 1 g water @ 25º C
 A couple of state functions for this sytem are
amount of mass (1 g) and T (25º C)
 There are others we will learn about
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Path independent
E.g., state would be the same if you condensed
steam or melted ice
 For the values of the state functions, it doesn’t
matter how the state got there
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Thermodynamic Laws
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Three laws – each derives a “new” state
function
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0th law: yields temperature (T)
1st law: yields enthalpy (H)
2nd law: yields entropy (S)
Zeroth law
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If two systems are in thermal equilibrium
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No heat is exchanged between the systems
They have the same “temperature”
T is the newly defined state function
How is temperature defined?
Measurement of T
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Centigrade
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100 divisions between melting and boiling
point of water
Kelvin - Based on Charles law
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At constant P and m, there is a linear
relationship between volume of gas and T
V = a1 + a2T
Where V = volume
T = temperature
a1 & a2 = constants
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Size of unit is same as centigrade
V (L)
Fig. Levine
T (ºC)
• 1 mole of N2 at constant P
• Experimental results:
- extrapolation of results show intercept T
@ V = 0 is about -273ºC
- Kelvin scale based on triple point of water
- defined as being 273.16 K
First law
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Change in the internal energy of a system
is the sum of the heat added (q) and
amount of work done (w) on system
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Energy conserved
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Three types of energy
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Kinetic and potential – physically defined
Internal – chemically defined
Three forms of energy
Potential + Kinetic energy
+ internal energy
Minimum or rest energy
Here only internal energy, U
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Internal energy (U)
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Molecular rotation, translation, vibration and
electrical energy
Potential energy of interactions of molecules
Relativistic rest-mass energy
In thermo, a system at rest
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Kinetic and potential energy = 0
Thermodynamics considers only changes in
internal energy
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New state function – Enthalpy (H)
H = U + PV
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PV = pressure * volume = work done
on/by the system
Units – energy, e.g. J, kJ, cal etc.
Extensive – i.e., additive.
Second Law
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A system cannot undergo a cyclic process
that extracts heat from a heat reservoir
and also performs an equivalent amount
of work on the surroundings
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i.e., it is impossible to build a machine that
converts heat to work with 100% efficiency
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New state function
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Entropy = S
Extensive = units of energy/T, e.g. kJ/K
Entropy is a variable used to defined Gibbs
free energy (G)
G used to determine equilibrium of
reactions
Equilibrium Thermodynamics
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Equilibrium occurs with a minimum of
energy in system
Systems not in equilibrium move toward
equilibrium through loss of energy
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If system is at constant T and P, measure
of energy of system is given by Gibbs free
energy (G)
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G = f(H,S,T)
G = H - TS
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G and H units = kJ (kcal)
S units = kJ/K (kcal/K)
T is Kelvin scale (K)
Imagine some system with A, B, C, and D components:
A+B↔C+D
Equilibrium A, B, C, and D present
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Consider processes in system at constant
T&P
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“Process” means system changes
May be chemical reaction
DG =DH - TDS
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Here D is change in state:
D = State2 – State1
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For all properties: G, H, T or S
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When system moves toward equilibrium:
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may release heat, e.g. DH < 0
entropy may increase, e.g. DS > 0
Both may happen
Thus:
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DG < 0 for spontaneous reaction
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G2 < G1; DG = G2 – G1 < 0
DG = 0 for process at equilibrium
Possible to calculate DG, and thus determine
(1) if reaction will occur spontaneously, and
(2) which way reaction will go.
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Non-equilibrium system:
A+B→C+D
DG ≠ 0
A+B←C+D
DG ≠ 0
Equilibrium system
A+B↔C+D
DG = 0
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G is an extensive state variable
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The amount of G in a system is divided
among components
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Need to know how G changes for each
component
First look at what variables control G
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It depends on the amount of material
What is G a function of?
Want to know how G changes if all (or any)
other variable change
Change = calculus
Math Review
(on board)
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If system is in thermal and mechanical
equilibrium:
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G = f(P, T, n1, n2, n3…)
Then total differential:
(on board)
Infinitesimal change in G caused by
infinitesimal change in P, T, n1, n2, n3…
These are values we need to know to
know DG
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Last term defined by Gibbs as chemical
potential (m)
(on board)
m is the amount that G changes (per mole)
with addition of new component
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Intensive property (G extensive)
Doesn’t depend on mass of system
For one component system m = G/n
For system at equilibrium, m of all
components are identical
Equilibrium, activities, chemical
potentials
(on board)