Lecture Notes Statistical Mechanics Alejandro L. Garcia San Jose State University PHYSICS 260 / Spring 2012 ⃝c2012 by Alejandro L. Garcia Department of Physics, San Jose State University, San Jose CA 95192-0106 Creative Commons Attribution- Noncommercial-Share Alike 3.0 United States License Preface Theselecturenotesshouldsupplementyournotes,especiallyforthoseofyou(likemyself)whoarepoor note takers. But I don’t recommend reading these notes while I lecture; when you go to see a play you don’t take the script to read along. Continue taking notes in class but relax and don’t worry about catching every little detail. WARNING: THESE NOTES ARE IN DRAFT FORM AND PROBABLY HAVE NUMEROUS TYPOS; USE AT YOUR OWN RISK. The main text for the course is R. K. Pathria and Paul D. Beale, Statistical Mechanics, 3rd Ed., Pergamon, Oxford (2011). Other texts that will sometimes be mentioned and that you may want to consult include: K. Huang, Statistical Mechanics, 2nd Ed., Wiley, New York (1987); F. Reif, Funda- mentals of Statistical and Thermal Physics, McGraw-Hill, New York (1965); L.D. Landau and E.M. Lifshitz, Statistical Physics, Addison-Wesley, Reading Mass. (1969). Here is the general plan for the course: First review some basic thermodynamics and applications; this is not covered in your text book but any undergraduate thermodynamics book should be suitable. Given the entropy of a system, all other thermodynamic quantities may be found. The theory of statisticalensembles(Chapters1−4inPathria)willtellushowtofindtheentropygiventhemicroscopic dynamics. The theory of ensembles connects mechanics and thermodynamics. The calculations involved in ensemble theory can be quite hard. We start with the easier non- interacting systems, such as paramagnets, classical ideal gas, Fermion gas, etc., (Chapters 6−8) and work up to interacting systems, such as condensing vapor and ferromagnets (Chapter 12−13). Most of statistical mechanics is concerned with calculating statistical averages but we briefly con- sider the variation of microscopic and macroscopic quantities about their mean values as predicted by fluctuation theory (Chapter 15). Finally we’ll touch on my own specialty, computer simulations in statistical mechanics in Chapter 16. Alejandro L. Garcia 1 Chapter 1 Thermodynamics Equation of State Lecture 1 We consider thermodynamic systems as physical systems entirely described by a set of thermodynamic parameters. Commonly used parameters include: • pressure, P, and volume, V • tension, τ and length, L • magnetic field, H and magnetization, M These parameters appear in pairs, the first being a generalized force and the second a generalized displacement, with the product of the pair having the dimensions of energy. All these parameters are well-defined by mechanics and are readily measured mechanically. An additional thermodynamic parameter, temperature, T, is not a mechanical parameter; we defer its definition for now. Rather than trying to be general, let’s say that our thermodynamic parameters are P, V, and T, for example, if our system was a simple gas in a container. Anisolatedsystem(noenergyormattercanenterorleavethesystem)will, intime, attainthermo- dynamicequilibrium. Atthermodynamicequilibrium,theparametersdonotchangewithtime. Though thermometers have been in use since the time of Galileo (late 1500’s) many concepts regarding temper- ature were not well understood until much later. In the mid 1700’s Joseph Black used thermometers to establish that two substances in thermal contact have the same at thermodynamic equilibrium. This is counterintuitive — if metal and wood are at thermal equilibrium the metal still “feels colder” than the wood. This misleading observation is due to the fact that our sense of touch uses conduction of heat and thus is a poor thermometer. The equation of state is the functional relation between P, V, and T at equilibrium (see Fig. 1.1). An example of an equation of state is the ideal gas law for a dilute gas, PV =NkT (Ideal gas only) where N is the number of molecules in the gas and k (k =1.38×10−23 J/K is Boltzmann’s constant. Boltzmann’sconstantdoesnothaveanydeepmeaning. Sincethetemperaturescalewasdevelopedlong before the relation between heat and energy was understood the value of k simply allows us to retain the old Kelvin and Celsius scales. Later we’ll see the theoretical justification for the ideal gas law in the next chapter. 2 CHAPTER 1. THERMODYNAMICS 3 Figure 1.1: Point in state space on the equation of state surface Figure 1.2: Path from points A to B in pressure-volume diagram You should not get the impression that the equation of state completely specifies all the thermody- namic properties of a system. As an example, both argon and nitrogen molecules have about the same mass and the two gases obey the ideal gas law yet their molar heat capacities are very different (can you guess why Ar and N are different?). 2 It follows from the equation of state that given V and T, we know P at equilibrium (or in general, given any two we know the third). First Law of Thermodynamics Call U(V,T) the internal energy∗ of our system (why don’t we write U(P,V,T)?). One way that the energy can change is if the system does mechanical work. The work done by a system in going from state A to B is, ∫ VB ∆W = P(V)dV AB VA since Work = Force × Distance = (Force/Area) × (Distance×Area) = (Pressure) × (Volume). Graph- ically, the work done is the area under the path travelled in going from A to B as shown in Fig. 1.2. The mental picture for this process is that of a piston lifting a weight as the system’s volume expands. Notice that if the path between A and B is changed, ∆W is different; for this reason dW =P dV is AB not an exact differential. ∗NOTATION:PathriawritesE(V,T)insteadofU(V,T). CHAPTER 1. THERMODYNAMICS 4 By conservation of energy, U(V ,T )−U(V ,T )=∆Q −∆W B B A A AB AB where∆Q isthenon-mechanicalenergychangeingoingfromAtoB. Wecall∆Q theheatadded AB AB to the system in going from A to B. If A and B are infinitesimally close, dU =dQ−dW =dQ−PdV (∗) where dU is an exact differential though dQ and dW are not. This is the first law of thermodynamics, which simply states that total energy is conserved. Theresultssofarareverygeneral. Ifourmechanicalvariablesaredifferent, everythingcarriesover. For example, for a stretched wire the mechanical variables are tension τ and length L instead of P and V. We simply replace P dV with −τdL in equation (∗). Note that the minus sign comes from the fact that the wire does work when it contracts (imagine the wire contracting and lifting a weight). Example: For n moles of an ideal gas held at constant temperature T find the work done in going 0 from volume V to V . A B Solution: ThenumberofmoleculesinnmolesisN =nN whereN =6.205×1023 isAvogadro’s A A number(nottobeconfusedwithAvocado’snumberwhichisthenumberofmoleculesinaguaca-mole). The ideal gas equation of state is PV =NkT =(nN )kT =nRT A where R=kN =8.31 J/(mol K) is the universal gas constant. A To find the work done, we solve, ∫ VB ∆W = P(V)dV AB VA Since the temperature is fixed at T , then P(V)=nRT /V or 0 0 ∫ VB dV ∆W =nRT =nRT ln(V /V ) AB 0 V 0 B A VA Note that on a different path we would have to know how temperature varied with volume, i.e., be given T(V) along the path. Heat capacity We define heat capacity, C, as dQ=CdT Specifically, the heat capacity at constant volume, C , is V ( ) ∂Q C = V ∂T V CHAPTER 1. THERMODYNAMICS 5 and the heat capacity at constant pressure ( ) ∂Q C = P ∂T P The heat capacity gives the amount of heat energy required to change the temperature of an object while holding certain mechanical variables fixed. We also define the heat capacity per unit mass, which is often called the specific heat, and the heat capacity per particle (e.g., c =C /N). V V For liquids and solids, since their expansion at constant pressure is often negligible, one finds C ≈ P C ; this is certainly not the case for gases. V For any function f(x,y) of two variables, the differential df is ( ) ( ) ∂f ∂f df = dx+ dy (Math identity) ∂x ∂y y x where the subscripts on the parentheses remind us that a variable is held fixed. Applying this identity to internal energy ( ) ( ) ∂U ∂U dU(T,V)= dT + dV (∗∗) ∂T ∂V V T If we hold volume fixed (i.e., set dV =0) in this expression and in (*) then we obtain the result ( ) ∂U dQ= dT (constant V) ∂T V From the above, ( ) ∂U C = V ∂T V Thisresultiseasytounderstandsincethechangeininternalenergyequalstheheataddedwhenvolume is fixed since no work can be done if volume is fixed. The heat capacity is important because given the equation of state and the heat capacity of a system, all other thermodynamic properties can be obtained. This is proved in most thermodynamics textbooks. Joule’s Free Expansion Experiment Let’s apply some of these results to a historically important experiment. Take a gas in an insulated container as shown in Fig. 1.3. Experimentally, one finds that T = T , that is, the temperature remains constant after the gas B A expands to fill both sides of the container. If you were thinking that the gas would cool when it expanded, I’ll explain the source of your confusion in a moment. In the free-expansion, the gas does no work so, ∆W = 0. The system is insulated so no heat AB enters or leaves, thus ∆Q =0. From first law, the internal energy U must remain constant. AB From (∗∗) for constant U, ( ) ( ) ∂U ∂U 0= dT + dV (dU =0) ∂T ∂V V T Since the experiment shows that the temperature does not change on expansion, then dT =0 and thus for an ideal gas, ( ) ∂U dV =0 (Ideal Gas) ∂V T CHAPTER 1. THERMODYNAMICS 6 Figure 1.3: Joule free expansion experiment This means that U does not depend on V so U(V,T) = U(T) for a dilute gas. This defines what one means by an ideal gas. The physical meaning of this result is that when the particles in an ideal gas move farther apart, the energy of the particles does not change. This means that there is zero potential energy associated with any interaction between the particles. In a real gas this is not exactly true since there is a weak interaction between the molecules. In an ideal gas we assume that the potential energy is negligible so all the energy is kinetic energy. Since the heat capacity at constant volume is defined as ( ) ∂U C = V ∂T V for an ideal gas, dU C = (Ideal gas) V dT or dU =C dT (Ideal gas) V If C constant then, V U(T)=C T +constant (Ideal gas) V IfwesayU(T =0)=0thenforanidealgasU(T)=C T. Laterwederivetheseresultsusingstatistical V mechanics. Note: When a gas expands in a piston, it does work and thus if the system is insulated, the gas cools (see Fig. 1.4). In this case ∆Q =0 as before but ∆W >0 so U <U and thus T <T . AB AB B A B A Example: FornmolesofanidealgasheldatconstanttemperatureT findtheheataddedingoing 0 from volume V to V . A B Solution: Since U(T) for an ideal gas, if T is fixed then so is the internal energy. The first law says, U −U =∆W −∆Q B A AB AB CHAPTER 1. THERMODYNAMICS 7 Figure 1.4: Illustration of an ideal gas cooling in an adiabatic expansion under pressure. since U =U , then ∆Q =∆W , so the heat added equals the work done for an isothermal path B A AB AB (see the previous example). Entropy Lecture 2 The first law of thermodynamics may be written as: dU =dQ−dW We have an expression for work in terms of thermodynamic parameters: ∫ VB dW =PdV or ∆W = PdV AB VA where∆W istheworkdonebythesystemingoingfromstate(T ,V )to(T ,V )alongsomegiven AB A A B B path. We now introduce a similar result for heat: ∫ SB dQ=TdS or ∆Q = TdS AB SA where ∆Q is the heat added to the system. Unfortunately, the entropy, S, is not as easily measured AB as other thermodynamic variables. We can write: ∫ dQ QB 1 dS = or S −S = dQ T B A T QA The differential dS is exact (like dU and dV) so if we select a path between A and B such that T =T is fixed, then: 0 ∫ 1 B 1 S −S = dQ= ∆Q B A T T AB 0 A 0 By experimentally measuring ∆Q , we obtain S. AB CHAPTER 1. THERMODYNAMICS 8 Figure 1.5: Carnot engine cycle in P −V diagram ButwhatifwecannotgetfromAtoB alonganisothermalpath? Alonganadiabaticpath(noheat added or removed) dQ = 0 so dS = 0. In general, we can build a path between any two points A and B by combining isothermal and adiabatic paths and thus calculate S −S . B A The entropy is important in that given S(U,V) for a physical system, all other thermodynamic properties may be computed. We derive and use this result when we develop statistical mechanics. Third Law of Thermodynamics The entropy is a function of the thermodynamic parameters so we may write it as S(T,V). The third law says: S(T =0,V)=0 that is, at zero Kelvin the entropy is zero. Thus, we can find the entropy for arbitrary T and V by computing: ∫ (T,V) dQ S(T,V)= T (T=0,V) along any convenient path. There are many deep meanings to the third law and I encourage you to plunge in and read about them on your own. Engines and Refrigerators A heat engine is described by a closed cycle in a P −V diagram for which the enclosed area (i.e., the workdone)ispositive(contourloopisclockwise). Iftheloopiscounterclockwisewehavearefrigerator. Call the heat added in a cycle Q and the work done W; we define engine efficiency as η =W/Q , + + that is, the fraction of the heat added that is converted into work. An important example is a Carnot engine. In a Carnot engine the cycle consists of a pair of isothermal paths and a pair of adiabatic paths (see Fig. 1.5). Notice that the Carnot cycle is a rectangle in the T −S diagram (see Fig. 1.6). The area enclosed in the P −V diagram is the work done. Since the cycle is a closed loop, our final and initial internal