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Physics project
Based on
Thermodynamics
Thermodynamics
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a system:
Some portion of the universe that you wish to study
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the surroundings:
The adjacent part of the universe outside the system
Changes in a system are associated with the transfer of energy
Natural systems tend toward states of minimum energy
Energy States
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Unstable:
falling or rolling
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Stable:
at rest in lowest energy state
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Metastable:
in low-energy perch
Figure 5-1. Stability states. Winter (2001) An
Introduction to Igneous and Metamorphic Petrology.
Prentice Hall.
History
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A brief history of thermodynamics begins with
Otto von Guericke who in 1650 built and designed
the world's first vacuum pump and created the
world's first ever vacuum (known as the
Magdeburg hemispheres). He was driven to make
a vacuum in order to disprove Aristotle's long-held
supposition that 'nature abhors a vacuum'. Shortly
thereafter, Irish physicist and chemist Robert
Boyle had learned of Guericke's designs and in
1656, in coordination with English scientist
Robert Hooke, built an air pump.
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Using this pump, Boyle and Hooke noticed the
pressure-temperature-volume correlation. In time,
Boyle's Law was formulated, which states that
pressure and volume are inversely proportional.
Then, in 1679, based on these concepts, an
associate of Boyle's named Denis Papin built a
bone digester, which was a closed vessel with a
tightly fitting lid that confined steam until a high
pressure was generated.
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Later designs implemented a steam release
valve that kept the machine from exploding.
By watching the valve rhythmically move
up and down, Papin conceived of the idea of
a piston and a cylinder engine. He did not,
however, follow through with his design.
Nevertheless, in 1697, based on Papin's
designs, engineer Thomas Savery built the
first engine.
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The term thermodynamics was coined by James
Joule in 1858 to designate the science of relations
between heat and power.[1] By 1849, "thermodynamics", as a functional term, was used in
William Thomson's paper An Account of Carnot's
Theory of the Motive Power of Heat.[9] The first
thermodynamic textbook was written in 1859 by
William Rankine, originally trained as a physicist
and a civil and mechanical engineering professor
at the University of Glasgow
Classical thermodynamics
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Classical thermodynamics is the original early 1800s variation of
thermodynamics concerned with thermodynamic states, and properties
as energy, work, and heat, and with the laws of thermodynamics, all
lacking an atomic interpretation. In precursory form, classical
thermodynamics derives from chemist Robert Boyle’s 1662 postulate
that the pressure P of a given quantity of gas varies inversely as its
volume V at constant temperature; i.e. in equation form: PV = k, a
constant. From here, a semblance of a thermo-science began to
develop with the construction of the first successful atmospheric steam
engines in England by Thomas Savery in 1697 and Thomas
Newcomen in 1712. The first and second laws of thermodynamics
emerged simultaneously in the 1850s, primarily out of the works of
William Rankine, Rudolf Clausius, and William Thomson (Lord
Kelvin).
Statistical thermodynamics
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With the development of atomic and molecular
theories in the late 19th century, thermodynamics
was given a molecular interpretation. This field is
called statistical thermodynamics, which can be
thought of as a bridge between macroscopic and
microscopic properties of systems.[11]
Essentially, statistical thermodynamics is an
approach to thermodynamics situated upon
statistical mechanics, which focuses on the
derivation of macroscopic results from first
principles.
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It can be opposed to its historical predecessor
phenomenological thermodynamics, which gives
scientific descriptions of phenomena with
avoidance of microscopic details. The statistical
approach is to derive all macroscopic properties
(temperature, volume, pressure, energy, entropy,
etc.) from the properties of moving constituent
particles and the interactions between them
(including quantum phenomena). It was found to
be very successful and thus is commonly used.
Chemical thermodynamics
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Chemical thermodynamics is the study of
the interrelation of heat with chemical
reactions or with a physical change of state
within the confines of the laws of
thermodynamics. During the years 1873-76
the American mathematical physicist Josiah
Willard Gibbs published a series of three
papers, the most famous being On the
Equilibrium of Heterogeneous Substances,
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in which he showed how thermodynamic
processes could be graphically analyzed, by
studying the energy, entropy, volume, temperature
and pressure of the thermodynamic system, in
such a manner to determine if a process would
occur spontaneously. During the early 20th
century, chemists such as Gilbert N. Lewis, Merle
Randall, and E. A. Guggenheim began to apply
the mathematical methods of Gibbs to the analysis
of chemical processes
Thermodynamic systems

An important concept in thermodynamics is the
“system”. Everything in the universe except the
system is known as surroundings. A system is the
region of the universe under study. A system is
separated from the remainder of the universe by a
boundary which may be imaginary or not, but
which by convention delimits a finite volume. The
possible exchanges of work, heat, or matter
between
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between the system and the surroundings
take place across this boundary. Boundaries
are of four types: fixed, moveable, real, and
imaginary.
Basically, the “boundary” is simply an
imaginary dotted line drawn around the
volume of a something in
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which there is going to be a change in the
internal energy of that something. Anything
that passes across the boundary that effects
a change in the internal energy of that
something needs to be accounted for in the
energy balance equation. That “something”
can be the volumetric region surrounding a
single atom
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resonating energy, such as Max Planck defined in
1900; it can be a body of steam or air in a steam
engine, such as Sadi Carnot defined in 1824; it can
be the body of a tropical cyclone, such as Kerry
Emanuel theorized in 1986 in the field of
atmospheric thermodynamics; it could also be just
one nuclide (i.e. a system of quarks) as some are
theorizing presently in quantum thermodynamics.

For an engine, a fixed boundary means the
piston is locked at its position; as such, a
constant volume process occurs. In that
same engine, a moveable boundary allows
the piston to move in and out. For closed
systems, boundaries are real while for open
system boundaries are often imaginary.
There are five dominant classes of systems:
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Isolated Systems – matter and energy may
not cross the boundary.
Adiabatic Systems – heat must not cross the
boundary.
Diathermic Systems - heat may cross
boundary.

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Closed Systems – matter may not cross the
boundary.
Open Systems – heat, work, and matter may
cross the boundary (often called a control
volume in this case).
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As time passes in an isolated system,
internal differences in the system tend to
even out and pressures and temperatures
tend to equalize, as do density differences.
A system in which all equalizing processes
have gone practically to completion, is
considered to be in a state of
thermodynamic equilibrium.

In thermodynamic equilibrium, a system's
properties are, by definition, unchanging in
time. Systems in equilibrium are much
simpler and easier to understand than
systems which are not in equilibrium.
Often, when analysing a thermodynamic
process,

it can be assumed that each intermediate
state in the process is at equilibrium. This
will also considerably simplify the situation.
Thermodynamic processes which develop
so slowly as to allow each intermediate step
to be an equilibrium state are said to be
reversible processes.
Thermodynamic states

When a system is at equilibrium under a
given set of conditions, it is said to be in a
definite state. The state of the system can be
described by a number of intensive
variables and extensive variables.

The properties of the system can be
described by an equation of state which
specifies the relationship between these
variables. State may be thought of as the
instantaneous quantitative description of a
system with a set number of variables held
constant.
Thermodynamic processes

A thermodynamic process may be defined
as the energetic evolution of a
thermodynamic system proceeding from an
initial state to a final state. Typically, each
thermodynamic process is distinguished
from other processes, in energetic character,
according to what parameters, as
temperature, pressure, or volume, etc., are
held fixed.
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The seven most common thermodynamic
processes are shown below:
An isobaric process occurs at constant pressure.
An isochoric process, or isometric/isovolumetric
process, occurs at constant volume.
An isothermal process occurs at a constant
temperature.
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An adiabatic process occurs without loss or gain
of heat.
An isentropic process (reversible adiabatic
process) occurs at a constant entropy.
An isenthalpic process occurs at a constant
enthalpy. Also known as a throttling process or
wire drawing.
A steady state process occurs without a change in
the internal energy of a system.
A review of basic thermodynamics: A refresher
The ball represents
mass exchange
The arrow represents
energy exchange
The First Law of Thermodynamics
• Heat and work are equivalent
• Energy is conserved in any transformation
• The change of energy of a system is independent of the path taken
Energy can be neither created nor detroyed
DE = q - w or dE = dq - dw
dE = dq - P dV
E = internal energy
q = heat
w = work
P = pressure
V = volume
Enthalpy
dE = dq - P dV
H = E + PV
dH = dq
H = enthalpy
The change in the enthalpy of a system (DH) during a reversible
change in state at constant pressure is equal to the heat absorbed
by the system during that change in state.
The enthalpy of formation of compounds and their ions and
molecules in aqueous solution is the heat absorbed or given off by
chemical reactions in which the compounds, ions, and molecules
form from the elements in the standard state (25°C, 1 atm)
Heats of Reaction
DH = S nH (products) - S nH (reactants)
n = molar coefficient of each reactant/product
When DH is positive, the reaction is endothermic (heat flows
from the surroundings to the system);
When DH is negative, the reaction is exothermic (heat flows
from the system to the surroundings
Heats of Reaction
DH = S nH (products) - S nH (reactants)
For example, evaporation: H2O(l)
H2O(g)
DH = H(H2O(g)) - H(H2O(l))
DH = (-57.80) - (-68.32) = 10.52 kcal
The reaction is endothermic (i.e., sweating is a mechanism
for cooling the body)
Heat Capacity
When heat is added to a solid, liquid, or gas, the temperature
of the substance increases:
dq = C dT
dq = dH
dH = C dT, at constant pressure (important in geochemistry)
C = heat capacity
T = temperature
Heat capacities vary with temperature…
The Second Law of Thermodynamics
• It is impossible to construct a machine that is able to convey heat
by a cyclical process from one reservoir at a lower temperature to
another at a higher temperature unless outside work is done
(i.e, air conditioning is never free)
• Heat cannot be entirely extracted from a body and turned into work
(i.e., an engine can never run 100% efficiently) — a certain fraction
of the enthalpy of a system is consumed by an increase in entropy
• Every system left to itself will, on average, change toward a
condition of maximum randomness — entropy of a system
increases spontaneously and energy must be spent to reverse this
tendency
The entropy of the universe always increases
or
“You can’t shovel manure into the rear end of a horse and
expect to get hay out of its mouth”
Entropy
DS = S nS (products) - S nS (reactants)
For example: H2O(l)
H2O(g)
DS = S(H2O(l)) - S(H2O(g))
DS = 45.10 - 16.71 = 28.39 cal/deg
When DS is positive, entropy of the system increases with
the change of state;
When DS is negative, entropy decreases
The fundamental equation of thermodynamics
The ratio of heat gained or lost to temperature will always
be the same, regardless of path, for a reversible reaction
dE = T dS - P dV
When dE = 0, T dS = P dV
dS = dq/T = P dV/T
dS = dq/T — reversible process
dS > dq/T — irreversible process
Look familiar?
dP DS
=
dT DV