Thermodynamics and Statistical Physics
When we consider work done we have to decide on a sign convention.
PH605 : Thermal and Statistical Physics 2
THERMODYNAMICS .....................................................................................................4
Review of Zeroth, First, Second and Third Laws...................................................4
Thermodynamics................................................................................................4
The zeroth law of thermodynamics,...................................................................4
Temperature, T...................................................................................................4
Heat, Q ...............................................................................................................4
Work, W.............................................................................................................4
PH605 Internal energy, U ..............................................................................................5
The first law of thermodynamics, ......................................................................5
Isothermal and Adiabatic Expansion .................................................................6
Heat Capacity.....................................................................................................6
Heat capacity at constant volume, CV ................................................................7
Heat capacity at constant pressure, CP ...............................................................7
Relationship between CV and CP .......................................................................8
Thermal and The second law of thermodynamics, .................................................................8
Heat Engines ......................................................................................................9
Efficiency of a heat engine ..............................................................................10
Statistical Physics The Carnot Cycle .............................................................................................11
The Otto Cycle .................................................................................................13
Concept of Entropy : relation to disorder............................................................15
The definition of Entropy.................................................................................16
M.J.D.Mallett Entropy related to heat capacity.......................................................................16
The entropy of a rubber band...........................................................................17
The third law of thermodynamics, ...................................................................18
P.Blümler The central equation of thermodynamics.........................................................18
The entropy of an ideal gas..............................................................................18
Thermodynamic Potentials : internal energy, enthalpy, Helmholtz and Gibbs
functions, chemical potential ...............................................................................19
Internal energy .................................................................................................20
Recommended text books: Enthalpy ...........................................................................................................20
Helmholtz free energy......................................................................................20
• Finn C.B.P. : Thermal Physics Gibbs free energy .............................................................................................21
• Adkins C.J. : Equilibrium Thermodynamics Useful work......................................................................................................21
• Mandl F: Statistical Physics Chemical Potential ...........................................................................................22
The state functions in terms of each other .......................................................22
Differential relationships : the Maxwell relations...............................................23
Maxwell relation from U .................................................................................23
Maxwell relation from H .................................................................................24
Maxwell relation from F ..................................................................................24
Maxwell relation from G .................................................................................25
Use of the Maxwell Relations..........................................................................26
Applications to simple systems.............................................................................26
The thermodynamic derivation of Stefan’s Law .............................................27
Equilibrium conditions : phase changes..............................................................28
Phase changes ..................................................................................................28
P-T Diagrams ...................................................................................................29
PVT Surface.....................................................................................................29
First-Order phase change .................................................................................30
Second-Order phase change.............................................................................31
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Phase change caused by ice skates...................................................................31
The Clausius-Clayperon Equation for 1st order phase changes. ......................32 Thermodynamics
The Ehrenfest equation for 2nd order phase changes .......................................33
BASIC STATISTICAL CONCEPTS ..................................................................................35
Isolated systems and the microcanonical ensemble : the Boltzmann-Planck Review of Zeroth, First, Second and Third Laws
Entropy formula ...................................................................................................35
Why do we need statistical physics ?...............................................................35
Macrostates and Microstates............................................................................35 Thermodynamics
Classical vs Quantum.......................................................................................36
The thermodynamic probability, Ω..................................................................36 Why study thermal and statistical physics ? What use is it ?
How many microstates ?..................................................................................36
What is an ensemble ?......................................................................................37
Stirling’s Approximation .................................................................................39
Entropy and probability.......................................................................................39 The zeroth law of thermodynamics,
The Boltzmann-Planck entropy formula..........................................................40
Entropy related to probability ..........................................................................40
If each of two systems is in thermal equilibrium with a third, then they are also in
The Schottky defect .........................................................................................41
thermal equilibrium with each other.
Spin half systems and paramagnetism in solids...................................................43
Systems in thermal equilibrium and the canonical ensemble : the Boltzmann
This implies the existence of a property called temperature. Two systems that
distribution...........................................................................................................45
are in thermal equilibrium with each other must have the same temperature.
The Boltzmann distribution .............................................................................45
Single particle partition function, Z, and ZN for localised particles : relation to
Helmholtz function and other thermodynamic parameters .................................47 Temperature, T
The single particle partition function, Z ..........................................................47
The partition function for localised particles ...................................................47 th
The 0 law of thermodynamics implies the existence of a property of a system
The N-particle partition function for distinguishable particles........................47
The N-particle partition function for indistinguishable particles.....................48 which we shall call temperature, T.
Helmholtz function ..........................................................................................49
Adiabatic cooling .............................................................................................50 Heat, Q
Thermodynamic parameters in terms of Z.......................................................53
In general terms this is an amount of energy that is supplied to or removed
from a system. When a system absorbs or rejects heat the state of the system
must change to accommodate it. This will lead to a change in one or more of
the thermodynamic parameters of the system e.g. the temperature, T, the
volume, V, the pressure, P, etc.
Work, W
When a system has work done on it, or if it does work itself, then there is a
flow of energy either into or out of the system. This will also lead to a change
in one or more of the thermodynamics parameters of the system in the same
way that gaining or losing heat, Q, will cause a change in the state of the
system, so too will a change in the work, W, done on or by the system.
When dealing with gases, the work done is usually related to a change in the
volume, dV, of the gas. This is particularly apparent in a machine such as a
cars engine.
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If a thermally isolated system is brought from one equilibrium state to another, the
work necessary to achieve this change is independent of the process used.
We can write this as,
dU = đ WAdiabatic
Note : when we consider work done we have to decide on a sign convention.
By convention, work done on a system (energy gain by the system) is positive
and work done by the system (loss of energy by the system) is negative.
e.g.
Internal energy, U
• đ W = + PdV : compression of gas in a pump (T of gas increases).
The internal energy of a system is a measure of the total energy of the • đ W = − PdV : expansion of gas in an engine (T of gas decreases).
system. If it were possible we could measure the position and velocity of every
particle of the system and calculate the total energy by summing up the
individual kinetic and potential energies. Isothermal and Adiabatic Expansion
N N When we consider a gas expanding, there are two ways in which this can
U = ∑ KE + ∑ PE occur, isothermally or adiabatically.
n =1 n =1
• Isothermal expansion : as it’s name implies this is when a gas
However, this is not possible, so we are never able to measure the internal expands or contracts at a constant temperature (‘iso’-same, ‘therm’-
energy of a system. What we can do is to measure a change in the internal temperature). This can only occur if heat is absorbed or rejected by the
energy by recording the amount of energy either entering or leaving a system. gas, respectively. The final and initial states of the system will be at the
same temperature.
In general, when studying thermodynamics, we are interested in changes of
state of a system.
∆U = ∆Q + ∆W
which we usually write,
dU = đ Q + đ W
The bar through the differential, đ , means that the differential is inexact, this
means that the differential is path dependent i.e. the actual value depends on
the route taken, not just the start and finish points.
• Adiabatic expansion : this is what happens when no heat is allowed
to enter or leave the system as it expands or contracts. The final and
initial states of the system will be at different temperatures.
Heat Capacity
As a system absorbs heat it changes its state (e.g. P,V,T) but different
systems behave individually as they absorb the same heat so there must be a
parameter governing the heat absorption, this is known as the heat capacity, C.
The first law of thermodynamics,
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The heat capacity of a material is defined as the limiting ration of the heat, Q, Relationship between CV and CP
absorbed, to the rise in temperature, ∆T, of the material. It is a measure of the
amount of heat required to increase the temperature of a system by a given The internal energy of a system can be written as,
amount.
dU = đ Q + đ W
Q ⇒ đ Q = dU - PdV
C = limit
∆T → 0 ∆T
Assuming the change of internal energy is a function of volume and
When a system absorbs heat its state changes to accommodate the increase temperature, U = U (V , T ) , i.e. we have a constant pressure process, this
of internal energy, therefore we have to consider how the heat capacity of a can be written as,
system is governed when there are restrictions placed upon how the system
can change.
∂U ∂U
đQ = dV + dT + PdV
In general we consider systems kept at constant volume and constant ∂V T ∂T V
temperature and investigate the heat capacities for these two cases.
which leads to,
Heat capacity at constant volume, CV
đ QP ∂U ∂V ∂U ∂V
⇒ CP = = + + P
If the volume of the system is kept fixed then no work is done and the heat dT ∂V T ∂T P ∂T V ∂T P
capacity can be written as,
∂U ∂V
∴CP = CV + P +
đQV ∂U ∂V T ∂T P
CV = =
dT ∂T V
This is the general relationship between CV and CP.
Heat capacity at constant pressure, CP In the case of an ideal gas the internal energy is independent of the volume
(there is zero interaction between gas particles), so the formula simplifies to,
The heat capacity at constant pressure is therefore analogously,
∂V
CP = CV + P
đQP ∂T P
CP =
dT ⇒ CP − CV = R
We now use a new state function known as enthalpy, H, (which we discuss The second law of thermodynamics,
later).
H = U + PV The Kelvin statement of the 2nd law can be written as,
⇒ dH = dU + PdV + VdP
It is impossible to construct a device that, operating in a cycle, will produce no
dH = đ Q + VdP effect other than the extraction of heat from a single body at a uniform temperature
and the performance of an equivalent amount of work.
Using this definition we can write,
đQP ∂H
CP = =
dT ∂T P
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It would be useful to convert all the heat , QH, extracted into useful work but
this is disallowed by the 2nd law of thermodynamics.
A more concise form of this statement is,
A process whose only effect is the complete conversion of heat into work is
impossible. If this process were possible it would be possible to join two heat engines
together, whose sole effect was the transport of heat from a cold reservoir to a
Another form of the 2nd law is known as the Clausius statement, hot reservoir.
It is impossible to construct a device that, operating in a cycle, will produce no
effect other than the transfer of heat from a colder to a hotter body.
Efficiency of a heat engine
Heat Engines
We can define the efficiency of a heat engine as the ratio of the work done to
the heat extracted from the hot reservoir.
Heat engines convert internal energy to mechanical energy. We can consider
taking heat QH from a hot reservoir at temperature TH and using it to do useful
W QH − QC Q
work W, whilst discarding heat QC to a cold reservoir TC. η= = = 1− C
QH QH QH
From the definition of the absolute temperature scale1, we have the
relationship,
QC QH
=
TC TH
1 For a proof of this see Finn CP, Thermal Physics,
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One way of demonstrating this result is the following. Consider two heat
engines which share a common heat reservoir. Engine 1 operates between T1
and T2 and engine 2 operates between T2 and T3. We can say that there must
be a relationship between the ratio of the heat extracted/absorbed to the
temperature difference between the two reservoirs, i.e.
Q Q Q
= f (θ ,θ
1
) , = f (θ ,θ
2 '
) , = f (θ ,θ
1 ''
)
Q Q Q
1 2 2 3 1 3
2 3 3
Therefore the overall heat engine can be considered as a combination of the
two individual engines.
f (θ ,θ
''
1 3
) = f (θ ,θ ) f (θ
1 2
'
2
,θ 3
)
However this can only be true if the functions factorize as,
The Carnot cycle is a closed cycle which extracts heat QH from a hot reservoir
and discards heat QC into a cold reservoir while doing useful work, W. The
T (θ x )
f (θ x ,θ y ) → cycle operates around the cycle A►B►C►D►A
T (θ y )
Where T(θ) represents a function of absolute, or thermodynamic temperature.
Therefore we have the relationship,
Q1 T (θ1 )
=
Q2 T (θ 2 )
Therefore we can also write the efficiency relation as,
TC
η = 1−
TH
The efficiency of a reversible heat engine depends upon the temperatures We can consider this cycle in terms of the expansion/contraction of an ideal
between which it operates. The efficiency is always PH605 : Thermal and Statistical Physics 13 PH605 : Thermal and Statistical Physics 14
A heat engine can also operate in reverse, extracting heat, QC from a cold The 4-stroke petrol engine follows the Otto cycle rather than the Carnot cycle.
reservoir and discarding heat, QH, into a hot reservoir by having work done on The actual cycle differs slightly from the idealised cycle to accommodate the
it, W, the total heat discarded into the hot reservoir is then, introduction of fresh petrol/air mixture and the evacuation of exhaust gases.
QH = QC + W
This is the principle of the refrigerator.
The Otto Cycle
The Carnot cycle represents the most efficient heat engine that we can
contrive. In reality it is unachievable.
Two of the most common heat engines are found in vehicles, the 4-stroke
petrol engine and the 4-stroke diesel engine.
The 4-stroke cycle can be considered as:
1. Induction : Petrol/Air mixture drawn into the engine cylinder.
2. Compression : Petrol/Air mixture compressed to a small volume by the
rising piston.
3. Power : Ignition of petrol/air mixture causes rapid expansion pushing
the piston down the cylinder
4. Exhaust : Exhaust gases evacuated from the cylinder by the rising
piston.
The Otto cycle and the Diesel cycle can be approximated by PV diagrams.
Otto cycle
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đQ
∫ R T
=0
The inequality of an irreversible process is a measure of the change of
entropy of the process.
final đQ
∆S = S −S =∫
T
final initial
initial
Diesel cycle
so for an infinitesimal part of the process we have,
đQ
dS ≥
T
The definition of Entropy
An entropy change in a system is defined as,
Concept of Entropy : relation to disorder đQ
dS =
T
We shall deal with the concept of entropy from both the thermodynamic and
the statistical mechanical aspects. The entropy of a thermally isolated system increases in any irreversible process and
is unaltered in a reversible process. This is the principle increasing entropy.
Suppose we have a reversible heat engine that absorbs heat Q1 from a hot
reservoir at a temperature T1 and discards heat Q2 into a cold reservoir at a The entropy of a system can be thought of as the inevitable loss of precision,
temperature T2, then from the efficiency relation we have, or order, going from one state to another. This has implications about the
direction of time.
Q1 Q2
= The forward direction of time is that in which entropy increases – so we can
T1 T2 always deduce whether time is evolving backwards or forwards.
nd
but from the 2 law we know that we cannot have a true reversible cycle,
there is always a heat loss, therefore we should rewrite this relationship as, Although entropy in the Universe as a whole is increasing, on a local scale it
can be decreased – that is we can produce systems that are more precise –
Q1 Q2 or more ordered than those that produced them. An example of this is creating
< a crystalline solid from amorphous components. The crystal is more ordered
T1 T2
and so has lower entropy than it’s precursors.
The heat absorbed in one complete cycle of the heat engine is therefore, On a larger scale – life itself is an example of the reduction of entropy. Living
organisms are more complex and more ordered than their constituent atoms.
đQ
∫ T
≤0
Entropy related to heat capacity
This is known as the Clausius inequality.
Suppose the heat capacity of a solid is CP=125.48 JK-1. What would be the
If we had a truly reversible heat engine then this would be, entropy change if the solid is heated from 273 K to 373 K ?
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Knowing the heat capacity of the solid and the rise in temperature we can The third law of thermodynamics,
easily calculate the heat input and therefore the entropy change.
đQ The entropy change in a process, between a pair of equilibrium states, associated
dS = with a change in the external parameters tends to zero as the temperature
T
approaches absolute zero.
We integrate over the temperature range to determine the total entropy
change. Or more succinctly,
T final dQ
∆S = S final − Sinitial = ∫ The entropy of a closed system always increases.
Tinitial T
T final
CP dT An alternative form of the 3rd law given by Planck is,
=∫
T T
initial
The entropy of all perfect crystals is the same at absolute zero and may be taken as
final
T −1
zero.
= CP ln = 39.2 JK
Tinitial In essence this is saying that at absolute zero there is only one possible state
for the system to exist in so there is no ambiguity about the possibility of it
existing in one of several different states.
The entropy of a rubber band
This concept becomes more evident when we consider the statistical concept
A rubber band is a collection of long chain polymer molecules. In its relaxed of entropy.
state the polymers are high disordered and entangled. The amount of disorder
is high and so the entropy of the system must be high. The central equation of thermodynamics
Stretching force The differential form of the first law of thermodynamics is,
dU = đ Q + đ W
Using our definition for entropy and assuming we are dealing with a
compressible fluid we can write this as,
dU = TdS - PdV
This is more usually written as,
If the rubber band is stretched then the polymers become less entangled and
align with the stretching force. They form a quasi-crystalline state. This is a TdS = dU + PdV
more ordered state and must therefore have a lower entropy.
This assumes that all the work done is due to changes of pressure, rather
The total entropy in the stretched state is made up of spatial and thermal
than changes of magnetisation etc.
terms.
STotal = S Spatial + SThermal The entropy of an ideal gas
If the tension in the band is rapidly reduced then we are performing an The specific heat capacity at constant volume for a gas is,
adiabatic (no heat flow) change on the system. The total entropy must remain
unchanged since there is no heat flow, but the spatial entropy has increased ∂U dU
so the thermal entropy must decrease this means the temperature of the CV = =
rubber band drops. ∂T V dT
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Substituting this into the central equation gives, The equilibrium conditions of a system are governed by the thermodynamic
potential functions. These potential functions tell us how the state of the
TdS = CV dT + PdV system will vary, given specific constraints.
If we consider one mole of an ideal gas and use lower case letters to refer to The differential forms of the potentials are exact because we are now dealing
molar quantities then we can write this as, with the state of the system.
Internal energy
RT
Tds = cV dT + dv
v This is the total internal energy of a system and can be considered to be the
dT dv sum of the kinetic and potential energies of all the constituent parts of the
ds = cv +R system.
T v ∞ ∞
Integrating both sides gives us, U = ∑ KE + ∑ PE
n =1 n =1
s = cv ln T + R ln v + s0
This quantity is poorly defined since we are unable to measure the individual
So the entropy of an ideal gas has three main terms, contributions of all the constituent parts of the system.
1. A temperature term – related to the motion, and therefore kinetic Using this definition of internal energy and the 2nd law of thermodynamics we
energy of the gas are able to combine the two together to give us one of the central equations of
2. A volume term – related to the positions of the gas particles thermodynamics,
3. A constant term – the intrinsic disorder term which is un-measurable.
TdS = dU + PdV
As an example of this can be used, consider gas inside a cylinder of volume,
V0. Suppose the volume of the cylinder is suddenly doubled. What is the This enables us to calculate changes to the internal energy of a system when
increase in entropy of the gas ? it undergoes a change of state.
Assuming this change occurs at constant temperature, we can write, dU = TdS − PdV
∆s = s2V0 − sV0 = R ln 2V0 − R ln V0 Enthalpy
2V This is sometimes erroneously called the heat content of a system. This is a
= R ln 0 = R ln 2 state function and is defined as,
V0
H = U + PV
If we were dealing with more than one mole of gas we could write this as,
We are more interested in the change of enthalpy, dH, which is a measure of
∆s = nR ln 2 the heat of reaction when a system changes state. In a mechanical system
= Nk B ln 2 this could be when we have a change in pressure or volume. In a
predominantly chemical system this could be due to the heat of reaction of a
change in the chemistry of the system.
Where n is the number of moles and N is the number of molecules. We will
return to this result when we look at the statistical definition of entropy.
dH = dU + PdV + VdP
Thermodynamic Potentials : internal energy, enthalpy, Helmholtz
Helmholtz free energy
and Gibbs functions, chemical potential
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The Helmholtz free energy of a system is the maximum amount of work However, some of the work done is useless work, suppose the actual gas
obtainable in which there is no change in temperature. This is a state function expansion was a by product of some chemical reaction, then we would want
and is defined as, to know how much useful work has actually been done.
F = U − TS WUseful = WTotal − WUseless
The change of Helmholtz free energy is given by, = WTotal − P0 ∆V
= −∆F − P0 ∆V
dF = dU − TdS − SdT
= −∆ ( F + PV
0 )
= − PdV − SdT
WUseful = −∆G
Gibbs free energy Therefore, the decrease in the Gibbs free energy tells us how much useful
work was done by this process.
The Gibbs free energy of a system is the maximum amount of work obtainable
in which there is no change in volume. This is a state function and is defined Chemical Potential
as,
G = H − TS This is important when the quantity of matter is not fixed (e.g. we are dealing
The change of Gibbs free energy is given by, with a changing number of atoms within a system). When this happens we
have to modify our thermodynamic relations to take account of this.
dG = dH − TdS − SdT
= VdP − SdT dU = TdS − PdV + µ dN
dF = − PdV − SdT + µ dN
It is obvious that the Helmholtz and Gibbs free energies are related,
dG = VdP − SdT + µ dN
∆G = ∆F + ∆ ( PV )
This means that there are several ways of writing the chemical potential, µ.
and the correct one to use has to be ascertained for the system in hand. For
example, a metal undergoes very small volume changes so we could use the ∂U ∂F ∂G
Gibbs function whereas a gas usually has large volume changes associated µ = = =
with it and we have to chose the function depending upon the situation.
∂N S V ∂N V T ∂N T P
, , ,
We can also show that the chemical potential can be written,
Useful work
G
µ=
Suppose we have a system that does work and that part of that work involves N
a volume change. If the system returns to its initial state of pressure and
temperature at the end of it doing some work then there is no temperature The chemical potential, µ, is the Gibbs free energy per particle, provided only one
change, i.e. type of particle is present.
• Initial temperature and pressure = T0 and P0
• Final temperature and pressure = T0 and P0 The state functions in terms of each other
Then because there is no overall temperature change, the maximum amount We can write infinitesimal state functions for the internal energy, U, the
of work done by the system is given by the decrease in the Helmholtz free enthalpy, H, the Helmholtz free energy, F and the Gibbs free energy, G.
energy, F, of the system.
dU = TdS − PdV
WTotal = −∆F dH = TdS + VdP
dF = − SdT − PdV
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dG = − SdT + VdP ∂U ∂U
dU = dS + dV
By inspection of these equations it would appear that there are natural ∂S V ∂V S
variables which govern each of the state functions.
which would then mean that we can write,
For instance, from the formula for the Helmholtz free energy we can assume
its natural variables are temperature and volume and therefore we can write, ∂U ∂U
T = and P = −
∂S V ∂V S
∂F ∂F
= −S and = −P
∂T V ∂V T moreover we can then write,
and from the formula for the Gibbs free energy, assuming its natural variables ∂P ∂T
are temperature and pressure, we have, = −
∂S V ∂V S
∂G ∂G
= −S and =V This is the first Maxwell relation.
∂T P ∂P T
This means that if we know one of the thermodynamic potentials in terms of Maxwell relation from H
its natural variables then we can calculate the other state functions from it.
We already have an equation of state for dH,
Suppose we know the Gibbs free energy, G, in terms of its natural variables T
and P, then we can write, dH = TdS + VdP
∂G ∂G This suggests that H is a function of S and P, therefore we could rewrite this
U = G − PV + TS = G − P −T as,
∂P T ∂T P
∂G ∂H ∂H
H = G + TS = G − T dH = dS + dP
∂T P ∂S P ∂P S
∂G
F = G − PV = G − P which would then mean that we can write,
∂P T
∂H ∂H
T = and V =
Differential relationships : the Maxwell relations ∂S P ∂P S
moreover we can then write,
The Maxwell relations are a series of equations which we can derive from the
equations of state for U, H, F and G.
∂V ∂T
Maxwell relation from U =
∂S P ∂P S
We already have an equation of state for dU, This is the second Maxwell relation.
dU = TdS − PdV
Maxwell relation from F
This suggests that U is a function of S and V, therefore we could rewrite this
as, We already have an equation of state for dF,
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dF = − SdT − PdV
This suggests that F is a function of T and V, therefore we could rewrite this Use of the Maxwell Relations
as,
Consider applying pressure to a solid (very small volume change), reversibly
∂F ∂F and isothermally. The pressure applied changes from P1 to P2 at a
dF = dT + dV temperature, T. The process is reversible so we can write,
∂T V ∂V T
đQR
which would then mean that we can write, dS =
T
∂F ∂F
S = − and P = − Since the only variables we have are pressure, P and temperature, T, we can
∂T V ∂V T write the entropy change of the system as a function of these two variables.
moreover we can then write, ∂S ∂S
dS = dP + dT
∂P T ∂T P
∂P ∂S
= − so therefore,
∂T V ∂V T
∂S
This is the third Maxwell relation. dQR = TdS = T dP
∂P T
Maxwell relation from G The second term is zero since the process is isothermal, dT=0.
We already have an equation of state for dG, Using the Maxwell relation derived from the Gibbs free energy,
dG = VdP − SdT ∂V ∂S
= −
∂T P ∂P T
This suggests that G is a function of P and T, therefore we could rewrite this
as,
we can then write,
∂G ∂G
dG = dP + dT ∂V
∂P T ∂T P dQR = −T dP = −TV β dP
∂T P
which would then mean that we can write, Where β is the coefficient of expansion. Integrating this gives,
∂G ∂G P2
V = and S = − Q = − ∫ TV β dP
∂P T ∂T P P1
≈ −TV β ( P2 − P1 )
moreover we can then write,
The approximation sign assumes β to be constant.
∂V ∂S
= − Applications to simple systems
∂T P ∂P T
This is the fourth Maxwell relation.
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The thermodynamic derivation of Stefan’s Law 1 du 1
u= T − u
3 dT 3
The energy density of thermal radiation (black-body radiation) is dependent
4 T du
only on the temperature of the black body. The energy density at a particular u=
wavelength is then, 3 3 dT
uλ=uλ(λ,T) dT du
4 =
and the total energy density is, T u
u=u(T) ⇒ u = AT 4
where A is a constant.
The Planck formula for the spectral energy density is given by,
β 1
uλ ( λ , T ) = 5
λ e λkBT − 1
hc
where β is a constant.
The perfect gas theory tells us that the pressure of a gas can be expressed in
terms of the mean velocity, c.
P = 1 ρ c2
3
but when we are dealing with thermal radiation we can use the Einstein mass-
energy relation,
E = mc 2 ⇒ u = ρ c2
therefore,
u
P=
3
The energy equation for a PVT system is,
∂U ∂P
=T −P
∂V T ∂T V Equilibrium conditions : phase changes
substituting into this we have,
Phase changes
A change of phase of a system occurs when the system changes from one
distinct state into another. This change of phase can be caused by many
different factors e.g. temperature changes can cause a phase change
between a solid and a liquid, applied magnetic fields can cause a phase
change between a superconductor and a normal conductor.
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P-T Diagrams
In a simple system comprising a single substance we can construct a
Pressure-Temperature diagram (P-T diagram) showing how changes in
pressure or temperature can affect the phase of the system.
Both the PT and PV diagrams can be obtained from the PVT surface by
projection.
PVT Surface
The PT diagram is a specialised case of the more general PVT surface. This
gives us all the thermodynamic information we require when determining the
available phase changes.
First-Order phase change
A first-order phase change of a substance is characterised by a change in the
specific volume between the two phase, accompanied by a latent heat.
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at constant temperature and the phase of the ice crosses the melting line on
Some typical examples of first-order phase changes are: the PT diagram. The skater is now standing on a thin layer of water which acts
as a lubricant between the skate and the bulk of the ice.
• The transition of a solid melting into a liquid.
• The transition of a liquid boiling into a gas. If the temperature is too low then the increased pressure of the skate on the
• The change from superconductor to normal conductor, provided the ice surface is insufficient to cause the ice to cross the melting line and so ice
change occurs in an applied magnetic field. skating is not possible.
Skiing is not a pressure-melting effect since the surface area of the sky is far
Second-Order phase change to large to cause melting of snow, instead the effect is caused by frictional
heating and wax lubricant applied to the ski.
Second-Order phase change of a substance is characterised by no change in the
specific volume between the two phases and no accompanying latent heat.
The Clausius-Clayperon Equation for 1st order phase changes.
Some typical examples of a second-order phase change are:
When a phase change occurs we are mostly interested in how it affects the
• The transition from ferro-magnet to para-magnet at the Curie Gibbs free energy. At thermodynamic equilibrium the Gibbs function is at a
temperature. minimum, so at the transition line on the PT diagram the specific Gibbs energy
• The transition from superconductor to normal conductor, provided there is the same for both phases.
is no applied magnetic field.
• The change from normal liquid 4Helium to superfluid liquid 4Helium g1 ( P, T ) = g 2 ( P, T )
below 2.2K
Further along the transition line we must also have the same condition.
Phase change caused by ice skates g1 ( P + dP, T + dT ) = g 2 ( P + dP, T + dT )
Water has a particularly interesting phase diagram. It is one the few materials Expanding this to first order using a Taylor approximation,
that expands on freezing (or contracts on melting), this is a consequence of
the effects of Hydrogen-bonding within the material.
∂g ∂g ∂g ∂g
g1 ( P, T ) + 1 dT + 1 dP = g2 ( P, T ) + 2 dT + 2 dP
∂T P ∂P T ∂T P ∂P T
So therefore,
∂g1 ∂g 2 ∂g2 ∂g1
− dT = − dP
∂T P ∂ T P ∂P T ∂P T
From our previous discussion about the natural variables for the Gibbs
function, we can write,
∂G ∂G
V = and S = −
∂P T ∂T P
and therefore considering the specific quantities, we have
This has implications which can be used to our advantage to allow ice skating.
The ‘hollow-ground’ edge of an ice skate causes an enormous pressure on
the ice surface of 100 atmospheres or more. This pressure increases occurs
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dP s2 − s1 S2 − S1 This is the first Ehrenfest equation for a 2nd order phase change, the second
= = Ehrenfest equation can be derived by considering the continuity of the volume
dT v2 − v1 V2 − V1 of the two phases, which gives,
The latent heat associated with this phase change is related to the entropy dP β 2 − β1
change, =
dT κ 2 − κ1
L = ∆Q = T ( S2 − S1 )
where κ is the bulk modulus of the phase.
Which gives us the Clausius-Clayperon equation which is the gradient of the
phase transition line for a first order phase change.
dP L
=
dT T (V2 − V1 )
The Ehrenfest equation for 2nd order phase changes
In a second order phase change there is no latent heat, and so no entropy
change, and no volume change. So using a similar argument as the one for
the Clausius-Clayperon equation,
s1 ( P, T ) = s2 ( P, T )
Further along the transition line we must also have the same condition.
s1 ( P + dP, T + dT ) = s2 ( P + dP, T + dT )
Expanding this to first order using a Taylor approximation,
∂s ∂s ∂s ∂s
s1 ( P, T ) + 1 dT + 1 dP = s2 ( P, T ) + 2 dT + 2 dP
∂T P ∂P T ∂T P ∂P T
Now substituting in expressions for the specific heat capacity and the volume
expansion,
∂s 1 ∂v
c=T and β=
∂T P v ∂T P
we can write,
dP CP1 − CP 2
=
dT TV ( β1 − β 2 )
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Classical vs Quantum
Basic statistical concepts
The classical view of the universe allows us an infinite choice of position and
velocity. This implies that for any given macrostate there must be an infinite
Isolated systems and the microcanonical ensemble : the number of microstates that correspond to it.
Boltzmann-Planck Entropy formula
The quantum viewpoint however tells us that at the atomic level we are able to
assign quantum numbers to the properties of particles. So for a closed system
Why do we need statistical physics ? there are only a finite number of states that the system can occupy.
There are two ways that we can look at the behaviour of matter In this scenario the most probable macrostate for a system is the one that has
the most microstates – simply because we have the greatest chance of finding
1. Mechanical viewpoint : measuring the positions and velocities of atoms the system in that microstate.
at the microscopic level.
2. Thermodynamic viewpoint : measuring the bulk properties of matter at
The thermodynamic probability, Ω
the macroscopic level.
However, when we try and reconcile the two viewpoints we run into a The number of possible microstates for a given macrostate is called the
problem. Theoretically we should get the same answers from both viewpoints. thermodynamic probability, Ω, or sometimes, W. This is not a probability in the
However there is a problem with the direction of time, normal sense since it has a value greater than or equal to one.
• Microscopic time : reversible e.g. the laws of motion look identical Consider a system of two particles A and B that can both exist in one of two
whichever way time goes. energy levels, E1 and E2. The macrostate of this system can be defined by the
• Macroscopic time : non-reversible, there is a preferred direction of time total energy of the system.
defined by the increase of entropy of the Universe.
Macrostate (1) E1 + E1 (2) E1 + E2 (3) E2 + E2
The two viewpoints can be reconciled by looking at systems from a statistical A(E1),B(E2)
Microstate A(E1),B(E1) A(E2),B(E2)
or probabilistic viewpoint. A(E2),B(E1)
Macrostates and Microstates Ω 1 2 1
When we are dealing with a system, e.g. a gas inside a container, we can Therefore if both energy levels, E1 and E2 are equally likely the system has a
determine various properties of the system. In our example of a gas we can 50% chance of being in macrostate (2) and a 25% chance each of being in
determine, i.e. measure, the volume, V, the pressure, P, the temperature, T, macrostates (1) and (3).
the molecular density, ρ, plus others.
However, in general not every energy level is equally likely so the most likely
The state of the system can therefore be defined by quantities such as V,P,T macrostate is also governed by the probability of energy level occupation.
& ρ. This is called a macrostate. In thermodynamic terms we usually only
consider macrostates. This leads on to the concept of the partition function, Z, for a system, which
we will cover later.
However, this gives us no knowledge of the properties of the individual gas
particles, for instance their positions or velocities. There are several, How many microstates ?
effectively infinitely many, states of a system which all have the same
macrostate. Each of these individual states are known as microstates. They
have different positions, velocities etc. for the individual particles but all have Suppose we are dealing with a paramagnetic solid. This means we have
the same thermodynamic properties such as volume, pressure and atoms arranged in a regular crystalline lattice and each atom acts as a
temperature. magnetic dipole because it has an unpaired electron.
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In a magnetic field the magnetic dipole can either point along the field or
against the field. Therefore there are two possible states for each dipole,
‘spin-up’ or ‘spin-down’.
Therefore, for N atoms (N~1023) there must be 2N possible microstates for the
magnetic dipoles.
How long would it take for the system to sample all the possible microstates ?
Assuming a dipole changes orientation every 10-10seconds there would be a
new microstate generated every 10-33seconds, so every microstate would be
sampled after
~ 10
(223 − 33) seconds
The lifetime of the Universe is only ~1017seconds !
Instead of considering one system and watching it change we use the concept
of an ensemble.
What is an ensemble ?
An ensemble is a very large number of replica systems with identical
specifications e.g. volume, temperature, chemistry, number of particles etc.
The ensemble represents the macrostate of the system while each individual
replica represents on of the possible microstates.
System Ensemble
i 1 2 3 n
There are several different ensembles that we might encounter. The type of
ensemble is governed by the measurable parameters.
1. Micro-canonical ensemble : isolated systems, the total internal
energy, U, and number of particles, N, is well defined.
2. Canonical ensemble : systems in thermal equilibrium, the
temperature, T, and number of particles, N, is well defined.
3. Grand canonical ensemble : systems in thermal and chemical
contact, the temperature, T, and chemical potential, µ, is well defined.
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Stirling’s Approximation Alternatively you can stick with the 6-dimensional phase space and have N
points in it. In this latter picture you might cut up the phase space into cells
We are now dealing with very large numbers ! Almost any macroscopic and count how many particles are in each cell.
system will have on the order of 1023 particles.
From this picture of phase space cell occupation numbers it is a small step to
The mathematics of such large numbers can become very involved but the thermodynamic probability. If the total number of particles is N and the
fortunately there are some simple approximations that work well when the N is number of particles in each cell is N , then the thermodynamic probability is
j
very large. given by,
There are two particularly useful approximations known as Stirling’s two-term N!
and three-term approximations; they are, Ω= ≈ exp N ln N − ∑ N ln N
∏N !
j j
j j
• Stirling’s 2-term approximation j
The quantity, Ω, is known as the thermodynamic probability or thermodynamic
ln ( N !) = N ln N − N
weight of the state of an ensemble (Note: Ω is sometimes written as W). The
states with the largest values of Ω are those that are most likely to occur.
Example
ln (100!) ≈ 100ln100 − 100 However the states with the largest value of Ω are also those with the most
disorder – simply because there are so many configurations of the microstates
≈ 360.5170 ( Stirlings approximation )
to give one macrostate. This means that there must be relationship between
= 360.7394 ( Actual value) the thermodynamic weight of a system and the entropy of a system.
• Stirling’s 3-term approximation
Therefore if
Ω tends to a maximum,
ln ( N !) = N ln N − N + ln 2π N then,
S tends to a maximum.
Example
ln (100!) ≈ 100ln100 − 100 − ln 200π The Boltzmann-Planck entropy formula
≈ 360.7385 ( Stirlings approximation )
= 360.7394 ( Actual value) The Boltzmann-Planck equation for entropy is written,
These approximations are particularly useful when dealing with
S = k B ln Ω
thermodynamic probabilities.
where Ω is the thermodynamic probability, or the
number of arrangements, of the state of the system.
Entropy and probability The states with the largest value of Ω will be the ones
most likely to occur. This equation is carved on
In statistical mechanics each particle is seen as having its own dynamic state, Boltzmann’s tombstone in Vienna.
a position in space, and a spatial velocity or momentum. In three-dimensional
space this gives the particle 3 degrees of freedom.
Three position coordinates and three momentum coordinates place each Entropy related to probability
particle somewhere in the six-dimensional phase space. If there is more than
one particle you can consider the system of those particles as having 6 times To derive this we first consider an ensemble of a large number, v, of replica
N coordinates and hence a single position in a 6N-dimensional phase space. systems. Each of these individual systems can exist in one of, r, microstates,
and for each microstate there is an associated probability, pr, of the system
being in that microstate.
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