thermal physics: 425/625 acknowledgement: i acknowledge the excellent efforts by rutgers physics...
TRANSCRIPT
Thermal Physics: 425/625
• Acknowledgement: I acknowledge the excellent efforts by Rutgers Physics Professors Misha Gershenson and Weida Wu in preparing powerpoint course material for this textbook. The powerpoint presentations in this course are rooted in their work, with permission from Dr Wu, with modifications by me.
Physics 425/625 Thermal Physics Fall 2011
Professor: Pat Arnott
Office: Leifson Physics RM 213
e-mail: [email protected]
phone: 775-784-6834
office hour: Wed 1:00-3:00 PM and by appointment
Class: Tuesday and Thursday 11:00 am to 12:15 pm
Prerequisites: Physics 182, 301.
Physics 425/625 Thermal Physics Fall 2011
Textbook: An Introduction to Thermal Physics,
D.V. Schroeder, Addison-Wesley-Longman, 2000
Web Site Access: http://www.patarnott.com/phys625
Homework assignments, announcements and notes will be posted here. The link must be visited regularly.
Thermal Physics = Thermodynamics + Statistical Mechanics
- conceptually, the most important, enjoyable subject of the undergraduate physics program.
Thermodynamics provides a framework of relating the macroscopic properties of a system to one another. It is concerned only with macroscopic quantities and ignores the microscopic variables that characterize individual molecules (both strength and weakness).
Statistical Mechanics is the bridge between the microscopic and macroscopic worlds: it links the laws of thermodynamics to the statistical behavior of molecules.
Macroscopic Description is Qualitatively Different!
Why do we need to consider macroscopic bodies as a special class of physical objects?
For a single particle: all equations of classical mechanics, electromagnetism, and quantum mechanics are time-reversal invariant (Newton’s second law, F = dp/dt, looks the same if the time t is replaced by –t and the momentum p by –p).
For macroscopic objects: the processes are often irreversible (a time-reversed version of such a process never seems to occur). Examples: (a) living things grow old and die, but never get younger, (b) if we drop a basketball onto a floor, it will bounce several times and eventually come to rest - the arrow of time does exist.
“More is different”, Phil Anderson, Science, 177, 393 (1972)
Ancients struggled with heat: What is it?
http://www.infinite-energy.com/iemagazine/issue37/mysteries.html
“It is astonishing to realize that many modern conceptions (or “laws”) in the science of heat— thermodynamics— arose during the nineteenth century, a period of utter confusion about the fundamental nature of heat. How could it have been otherwise, given that the very existence of atoms was still in question! “
The Main Idea of the Course
Statistical descriptionof a large system
of identical (mostly, non-interacting) particles
Irreversibility of macro processes,
the 2nd Law of Thermodynamics
all microstates of an isolated system occur with the same probability, the concepts of multiplicity (configuration
space), Entropy
Equation of state for macrosystems
(how macroparameters of the system and the temperature
are interrelated)
Thermodynamic Systems, Macroscopic Parameters
Open systems can exchange both matter and energy with the environment.
Closed systems exchange energy but not matter with the environment.
Isolated systems can exchange neither energy nor matter with the environment.
Internal and external macroscopic parameters: temperature, volume, pressure, energy, electromagnetic fields, etc. (average values, fluctuations are ignored).
No matter what is the initial state of an isolated system, eventually it will reach the state of thermodynamic equilibrium (no macroscopic processes, only microscopic motion of molecules).
A very important macro-parameter: Temperature
Temperature is a property associated with random motion of many particles.
Introduction of the concept of temperature in thermodynamics is based on the the zeroth law of thermodynamics:
A well-defined quantity called temperature exists such that two systems will be in thermal equilibrium if both have the same temperature.
Temperature Measurement
Properties of a thermoscope (any device that quantifies temperature):
1. It should be based on an easily measured macroscopic quantity a (volume, resistance, etc.) of a common macroscopic system.
2. The function that relates the chosen parameter with temperature, T = f(a), should be monotonic.
3. The quantity should be measurable over as wide a range of T as possible.
The simplest case – linear dependence T = Aa (e.g., for the ideal gas thermometer, T = PV/NkB).
Thermometer a thermoscope calibrated to a standard temp. scale
T
a
the ideal gas thermometer, T = PV/NkB
the resistance thermometer with a semi- conductor sensor, TR exp
The Absolute (Kelvin) Temperature Scale
The absolute (Kelvin) temperature scale is based on fixing T of the triple point for water (a specific T =
273.16 K and P = 611.73 Pa where water can coexist in the solid, liquid, and gas phases in equilibrium).
TPP
PKT 16.273
- for an ideal gas constant-volume
thermoscope
absolute zero
T,K
PPTP
273.16
PTP – the pressure of the gas in a constant-volume gas thermoscope
at T = 273.16 K
0
Our first model of a many-particle system: the Ideal Gas
The ideal gas model - works well at low densities (diluted gases)
• all the molecules are identical, N is huge;
• the molecules are tiny compared to their average separation (point masses);
• the molecules do not interact with each other;
• the molecules obey Newton’s laws of motion, their motion is random;
• collisions between the molecules and the container walls are elastic.
Models of matter: gas models (random motion of particles)
lattice models (positions of particles are fixed)
Air at normal conditions:
~ 2.71019 molecules in 1 cm3 of air (Pr. 1.10)
Size of the molecules ~ (2-3)10-10 m, distance
between the molecules ~ 310-9 m
The average speed - 500 m/s
The mean free path - 10-7 m (0.1 micron)
The number of collisions in 1 second - 5 109
The Equation of State of Ideal Gases
P – pressure [Newtons/m2]
V – volume [m3]
n – number of moles of gas [mol]
T – the temperature in Kelvins [K]
R – a universal constant
nRTPV
8.315J
Rmol K
The ideal gas equation of state:
An equation of state - an equation that relates macroscopic variables (e.g., P, V, and T) for a given substance in thermodynamic equilibrium.
In equilibrium ( no macroscopic motion), just a few macroscopic parameters are required to describe the state of a system.
The Ideal Gas Law
isotherms
In terms of the total number of molecules, N = nNA
TNkPV Bthe Boltzmann constant kB = R/NA 1.3810-23 J/K
(introduced by Planck in 1899)
Avogadro’s Law: equal volumes of different gases at the same P and T contain the same amount of molecules.
The P-V diagram – the projection of the surface of the equation of state onto the P-V plane.
NA 6.022045×1023
Avogadro’s number
The equations of state cannot be derived within the frame of thermodynamics: they can be either considered as experimental observations, or “borrowed” from statistical mechanics.
Connection between KEtr and T for Ideal Gases
T of an ideal gas the kinetic energy of molecules
Pressure – the result of collisions between the molecules and walls of the container.
Strategy: Pressure = Force/Area = [p / t ]/Area
vx
Piston area A
L
Volume = LA
For each (elastic) collision: px = 2 m vx
Intervals between collisions: t = 2 L/vx
Momentum
Vmv
Vvp
AvL
mvP xxx
x
xi
111
/2
2 2
22x
N
ix vmNmvPV For N molecules -
?
no-relativisticmotion
Connection between KEtr and T for Ideal Gases (cont.)
22x
N
ix vmNmvPV
TNkPV B
22222
2
3
2
1
2
1xzyxtr vmvvvmvmKE
Tkvm Bx 2
TkKE Btr 2
3 - the temperature of a gas is a direct measure of the
average translational kinetic energy of its molecules!
Average kinetic energy of the translational motion of molecules:
TNkKEU Btr 2
3
The internal energy U of a monatomic ideal gas is independent of its volume, and depends only on T (U =0 for an isothermal process, T=const).
UPV3
2 - for an ideal gas of non-relativistic particles,
kin. energy (velocity)2 .
Comparison with Experiment
TNkU B2
3
- for a point mass with three degrees
of freedom
dU/dT(300K) (J/K·mole)
Monatomic
Helium 12.5
Argon 12.5
Neon 12.7
Krypton 12.3
Diatomic
H2 20.4
N2 20.8
O2 21.1
CO 21
Polyatomic
H20 27.0
CO2 28.5
Testable prediction: if we put a known dU into a sample of gas, and measure the resulting change dT, we expect to get
moleJ/K5.12
J/K1038.1mole1062
32
3
231-23
BNkdT
dU
Conclusion: diatomic and polyatomic gases can store thermal energy in forms other than the translational kinetic energy of the molecules.
Degrees of Freedom
The degrees of freedom of a system are a collection of independent variables required to characterize the system.
Diatomic molecules: 3 + 2 = 5 transl.+rotat. degrees of freedom
Polyatomic molecules: 6 (transl.+rotat.) degrees of freedom
Degrees of Freedom (cont.)
Plus all vibrational degrees of freedom. The one-dimensional vibrational motion counts as two degrees of freedom (kin. + pot. energies):
22
2
1
2
1xkxmxUK
For a diatomic molecule (e.g., H2), 5 transl.+rotat. degrees of freedom plus 2 vibrational degrees of freedom = total 7 degrees of freedom
Among 7 degrees of freedom, only 3 (translational) degrees correspond to a continuous energy spectrum (classical) , the other 4 – to a discrete energy
spectrum (Quantum).
x
U(x)
E1
E2
E3
E4
x
U(x)
-3
-2
-1
0
1
2
3
4
1.5 2.0 2.5 3.0 3.5 4.0
distance
Ene
rgy
Approx.
“Frozen” degrees of freedom
For an ideal gas
PV = NkBT
U = f/2 NkBT
U /kBT
3/2N
5/2N
7/2N
10 100 1000 T, K
Translation
Rotation
Vibration
one mole of H2
The rotational energy levels are ~15 meV apart, the difference between vibrational energy levels ~270 meV. Thus, the rotational degrees start contributing to U at T > 200 K, the vibrational degrees of freedom - at T > 3000 K.
Example of H2:
An energy available to a H2 molecule colliding with a wall at T=300 K: 3/2 kBT ~ 40 meV. If the difference between energy levels is >> kBT, then a typical collision cannot cause transitions to the higher (excited) states and thus cannot transfer energy to this degree of freedom: it is “frozen out”.
x
U(x)
E1
E2
E3
E4
kBT
Equipartition of Energy
vx
Piston – a mechanical system with one degree of freedom. Thus,
TkuMvm
Bx
2
1
22
22
M – the mass of a piston, u2 the average u2, where u is the piston’s speed along the x-axis.
Thus, the energy that corresponds to the one-dimensional translational motion of a macroscopic system is the same as for a molecule (in this respect, a macrosystem behaves as a giant “molecule”).
Equipartition Theorem: At temperature T, the average energy of any “quadratic” degree of freedom is 1/2kBT.
“Quadratic” degree of freedom – the corresponding energy = f(x2, vx2)
[ translational motion, (classical) rotational and vibrational motion, etc. ]
- holds only for a system of particles whose kinetic energy is a quadratic form of x2, vx
2 (e.g., the equipartition theorem does not work for photons, E = cp)
Questions to solve in class: Think about them beforehand.
a. What are the molecules in the air in this room? Give % estimates.
b. How many molecules are in the air in this room?
d. What is the total internal energy of the air in this room?
c. On average, how far apart are the molecules from each other?
Questions Continued
e. Assuming (bad assumption) that we could use all this energy to run a hair dryer (1 kW), how long could we run the air dryer on energy?
f. Assuming the walls, floor, and ceiling of the room are perfect blackbodies, how much power is emitted by them all?
Questions Continued
g. Use the idea of the deBroglie wavelength and the average kineticenergy to calculate the thermal deBroglie wavelength for the molecules in this room as a function of temperature. How does the thermal deBroglie wavelength compare with the inter particle spacingat room temperature? As T goes to zero? Interpret.
http://en.wikipedia.org/wiki/Thermal_de_Broglie_wavelength
http://en.wikipedia.org/wiki/Matter_wave
The root-mean-square speed
m
Tkvv B
rms
32 - not quite the average speed, but close...
For H2 molecules (m ~21.710-27 kg ) at 300K: vrms~ 1.84 103 m/s
For N2 – vrms (Pr. 1.18), for O2 – vrms= 461 m/s
v
D(v)
vrms
This speed is close to the speed of sound in the gas – the sound wave propagates due to the thermal motion of molecules.
Problem 1.16 The “exponential” atmosphere
Consider a horizontal slab of air whose thickness is dz. If this slab is at rest, the pressure holding it up from below must balance both the pressure from above and the weight of the slab. Use this fact to find an expression for the variation of pressure with altitude, in terms of the density of air, . Assume that the temperature of atmosphere is independent of height, the average molecular mass m.
z
z+dzP(z+dz)A
P(z)A Mg
area A AzPMgAdzzP )()(
A
MgzPdzzP )()(
AdzM gdz
dP
the density of air: Tk
Pm
Tk
PVN
V
Nm
V
M
BB
PkT
mg
dz
dP
kT
mgzPzP exp)0()(Assuming T is independent of z
The First Law of Thermodynamics (Ch.1)
Outline:
1. Internal Energy, Work, Heating
2. Energy Conservation – the First Law
3. Quasi-static processes
4. Enthalpy
5. Heat Capacity
Internal Energy
system U = kinetic + potential
system boundary
“environment”
The internal energy of a system of particles, U, is the sum of the kinetic energy in the reference frame in which the center of mass is at rest and the potential energy arising from the forces of the particles on each other.
Difference between the total energy and the internal energy?
The internal energy is a state function – it depends only on the values of macroparameters (the state of a system), not on the method of preparation of this state (the “path” in the macroparameter space is irrelevant).
U = U (V, T)In equilibrium [ f (P,V,T)=0 ] :
U depends on the kinetic energy of particles in a system and an average inter-particle distance (~ V-1/3) – interactions.
P
V T A
B
For an ideal gas (no interactions) : U = U (T) - “pure” kinetic
Internal Energy of an Ideal Gas
The internal energy of an ideal gas with f degrees of freedom: TNk
fU B2
f 3 (monatomic), 5 (diatomic), 6 (polyatomic)
How does the internal energy of air in this (not-air-tight) room change with T if the external P = const?
PVf
Tk
PVNTkN
fU
BroominBroomin 22
(here we consider only trans.+rotat. degrees of freedom, and neglect the vibrational ones that can be excited at very high temperatures)
- does not change at all, an increase of the kinetic energy of individual molecules with T is compensated by a decrease of their number.
Work and Heating (“Heat”)
We are often interested in U , not U. U is due to:
• Q - energy flow between a system and its environment due to T across a boundary and a finite thermal conductivity of the boundary
– heating (Q > 0) /cooling (Q < 0)(there is no such physical quantity as “heat”; to emphasize this fact, it is better to use the term “heating” rather than “heat”)
• W - any other kind of energy transfer across boundary - work
Heating/cooling processes:conduction: the energy transfer by molecular contact – fast-moving molecules transfer energy to slow-moving molecules by collisions;
convection: by macroscopic motion of gas or liquid
radiation: by emission/absorption of electromagnetic radiation.
HEATING
WORK
Work and Heating are both defined to describe energy transfer across a system boundary.
The First Law
For a cyclic process (Ui = Uf) Q = - W.
If, in addition, Q = 0 then W = 0
The first law of thermodynamics: the internal energy of a system can be
changed by doing work on it or by heating/cooling it.
U = Q + W conservation of energy.
P
V T
An equivalent formulation:
Perpetual motion machines of the first type do not exist.
Sign convention: we consider Q and W to be positive if energy flows into the system.
Quasi-Static Processes
Quasi-static (quasi-equilibrium) processes – sufficiently slow processes, any intermediate state can be considered as an equilibrium state (the macroparamers are well-defined for all intermediate states).
Examples of quasi-equilibrium processes:
isochoric: V = const
isobaric: P = const
isothermal: T = const
adiabatic: Q = 0
For quasi-equilibrium processes, P, V, T are well-defined – the “path” between two states is a continuous lines in the P, V, T space.
P
V T
1
2
Advantage: the state of a system that participates in a quasi-equilibrium process can be described with the same (small) number of macro parameters as for a system in equilibrium (e.g., for an ideal gas in quasi-equilibrium processes, this could be T and P). By contrast, for non-equilibrium processes (e.g. turbulent flow of gas), we need a huge number of macro parameters.
Work
The sign: if the volume is decreased, W is positive (by compressing gas, we increase its internal energy); if the volume is increased, W is negative (the gas decreases its internal energy by doing some work on the environment).
2
1
),(21
V
VdVVTPW
The work done by an external force on a gas enclosed within a cylinder fitted with a piston:
W = (PA) dx = P (Adx) = - PdV
x
P
W = - PdV - applies to any shape of system boundary
The work is not necessarily associated with the volume changes – e.g., in the Joule’s experiments on determining the “mechanical equivalent of heat”, the system (water) was heated by stirring.
dU = Q – PdV
A – the piston area
force
W and Q are not State Functions
P
V
P2
P1
V1 V2
A B
CD
- the work is negative for the “clockwise” cycle; if the cyclic process were carried out in the reverse order (counterclockwise), the net work done on the gas would be positive.
01212
211122
VVPP
VVPVVPWWW CDABnet
2
1
),(21
V
VdVVTPW
- we can bring the system from state 1 to state 2 along infinite # of paths, and for each path P(T,V) will be different.
U is a state function, W - is not thus, Q is not a state function either.
U = Q + W
Since the work done on a system depends not only on the initial and final states, but also on the intermediate states, it is not a state function.
PV diagram
P
V T
1
2
the difference between the values of some (state) function z(x,y) at these points:
Comment on State Functions
U, P, T, and V are the state functions, Q and W are not. Specifying an initial and final states of a system does not fix the values of Q and W, we need to know the whole process (the intermediate states). Analogy: in classical mechanics, if a force is not conservative (e.g., friction), the initial and final positions do not determine the work, the entire path must be specified.
x
y z(x1,y1)
z(x2,y2)
dyy
zdx
x
zzd
xy
, ,x ydz A x y dx A x y dy - it is an exact differential if it is
x
yxA
y
yxA yx
,,
, ,dz z x dx y dy z x y
A necessary and sufficient condition for this:
If this conditionholds:
y
yxzyxA
x
yxzyxA yx
,
,,
,
e.g., for an ideal gas:
dV
V
TdT
fNkPdVdUQ B 2
- cross derivativesare not equal
dVPSdTUd
U
V S- an exact differential
In math terms, Q and W are not exact differentials of some functions of macroparameters. To emphasize that W and Q are NOT the state functions, we will use sometimes the curled symbols (instead of d) for their increments (Q and W).
Problem
Imagine that an ideal monatomic gas is taken from its initial state A to state B by an isothermal process, from B to C by an isobaric process, and from C back to its initial state A by an isochoric process. Fill in the signs of Q, W, and U for each step.
V, m3
P, 105 Pa
A
BC
Step Q W U
A B
B C
C A
2
1
1 2
+ -- 0
-- + --
+ 0 +
T=const
TNkf
U B2 BPV Nk T
The Enthalpy
Isobaric processes (P = const):
dU = Q - PV = Q -(PV) Q = U + (PV)
The enthalpy is a state function, because U, P, and V are state functions. In isobaric processes, the energy received by a system by heating equals to the change in enthalpy.
Q = H
isochoric:
isobaric:
in both cases, Q does not depend on the path from 1 to 2.
Consequence: the energy released (absorbed) in chemical reactions at constant volume (pressure) depends only on the initial and final states of a system.
H U + PV - the enthalpy
The enthalpy of an ideal gas:(depends on T only)
TNkf
TNkTNkf
PVUH BBB
1
22
Q = U
Heat Capacity
T
QC
The heat capacity of a system - the amount of energy
transfer due to heating required to produce a unit temperature rise in that system
C is NOT a state function (since Q is not a state function) – it depends on the path between two states of a system
T
V
T1
T1+dT
i
f1 f2 f3
The specific heat capacitym
Cc
( isothermic – C = , adiabatic – C = 0 )
Quasistatic Processes in an Ideal Gas
isochoric ( V = const )
isobaric ( P = const )
021 W
TCTTNkQ VB 02
31221
0),( 12
2
1
21 VVPdVTVPW
TCTTNkQ PB 02
51221
21QdU
2121 QWdU
V
P
V1,2
PV= NkBT1
PV= NkBT21
2
V
P
V1
PV= NkBT1
PV= NkBT212
V2
(see the last slide)
Isothermal Process in an Ideal Gas
1
221 ln),(
2
1
2
1V
VTNk
V
dVTNkdVTVPW B
V
V
B
V
V
f
iBfi V
VTNkW ln
Wi-f > 0 if Vi >Vf (compression)
Wi-f < 0 if Vi <Vf (expansion)
isothermal ( T = const ) :
V
P
PV= NkBT
V1V2
W
2121 WQ
0dU
Adiabatic Process in an Ideal Gas
adiabatic (thermally isolated system)
PdVdTNkf
dUTNkf
U BB 22
( f – the # of “unfrozen” degrees of freedom )
dTNkVdPPdVTNkPV BB PVPdVf
VdPPdV 2
fP
dP
fV
dV 21,0
21
constVPPVP
P
V
V
111
1
lnln
The amount of work needed to change the state of a thermally isolated system depends only on the initial and final states and not on the intermediate states.
021 Q 21WdU
to calculate W1-2 , we need to know P (V,T)
for an adiabatic process
2
1
),(21
V
V
dVTVPW
011
P
P
V
V P
dP
V
dV
V
P
V1
PV= NkBT1
PV= NkBT21
2
V2
Adiabaticexponent
Adiabatic Process in an Ideal Gas (cont.)
V
P
V1
PV= NkBT1
PV= NkBT21
2
22 2
1 1 1
11 11 2 1 1
1 1 1 12 1
1( , )
1
1 1 1
1
VV V
V V V
PVW P V T dV dV PV V
V
PVV V
1+2/31.67 (monatomic), 1+2/5 =1.4 (diatomic), 1+2/6 1.33 (polyatomic)(again, neglecting the vibrational degrees of freedom)
constVPPV 11
An adiabata is “steeper” than an isotherma: in an adiabatic process, the work flowing out of the gas comes at the expense of its thermal energy its temperature will decrease.
V2
Prove 1 2 2 2 B
f fW PV Nk T U
Summary of quasi-static processes of ideal gas
Quasi-Static process
U Q WIdeal gas
law
isobaric (P=0)
isochoric (V=0)
0
isothermal (T=0)
0
adiabatic (Q=0)
0
fi
i f
VV
T T
2 2B
f fU Nk T P V
f iU U U
2
2
fP V
P V
2 2B
f fU Nk T P V
2
fP V
fi
i f
PP
T T
i i f fPV P Vln fB
i
VNk T
VW
i i f fPV P V 2 2B
f fU Nk T PV U
Problem
Imagine that we rapidly compress a sample of air whose initial pressure is 105 Pa and temperature is 220C (= 295 K) to a volume that is a quarter of its original volume (e.g., pumping bike’s tire). What is its final temperature?
2211
222
111
VPVP
TNkVP
TNkVP
B
B
1
2
1
2
12
1
1121
2
11
2
112 T
T
V
VT
T
VPTNk
V
VP
V
VPP B
constVTVT 122
111
KKKV
VTT 51474.12954295 4.0
1
2
112
For adiabatic processes:
Rapid compression – approx. adiabatic, no time for the energy exchange with the environment due to thermal conductivity
constTP /1also
- poor approx. for a bike pump, works better for diesel engines
Non-equilibrium Adiabatic Processes
- applies only to quasi-equilibrium processes !!! constTV 1
2. On the other hand, U = Q + W = 0
U ~ T T – unchanged
(agrees with experimental finding)
Contradiction – because approach
#1 cannot be justified – violent
expansion of gas is not a quasi-
static process. T must remain the
same.
constTV 11. V – increases
T – decreases (cooling)
Free expansion
CV and CP
dT
PdVdU
dT
QC
V = const
P = const
VV T
UC
the heat capacity at constant volume
the heat capacity at constant pressure
PP T
HC
To find CP and CV, we need f (P,V,T) = 0 and U = U (V,T)
nRf
Nkf
C BV 22
For an ideal gas TNkf
U B2
# of moles
12 B
fH Nk T
12P
fC nR
For one mole of a monatomic ideal gas:
RCRC PV 2
5
2
3
Another Problem
During the ascent of a meteorological helium-gas filled balloon, its volume increases from Vi = 1 m3 to Vf = 1.8 m3, and the pressure inside the balloon decreases from 1 bar (=105 N/m2) to 0.5 bar. Assume that the pressure changes linearly with volume between Vi and Vf.(a) If the initial T is 300K, what is the final T?(b) How much work is done by the gas in the balloon?(c) How much “heat” does the gas absorb, if any?
P
V
Pf
Pi
Vi Vf
K2701mbar1
1.8mbar5.0K300
3
3
ii
ffif
BB VP
VPTT
Nk
PVTTNkPV
(a)
(b) f
i
V
V
ON dVVPW )(
bar625.1bar/m625.0 3 VVP
(c) ONWQU
Jmbarmbarmbar 333 41066.04.05.08.05.0)( f
i
V
V
BY dVVPW
- work done on a system f
i
V
V
BY dVVPW )( - work done by a system
BYON WW
JJJ 445 105.41061.0105.112
3
2
3
BY
i
fiiONifBON W
T
TVPWTTNkWUQ