The First Law Discussion

THE FIRST LAW (DISCUSSION)
 Thermodynamics is the study of the transformations of energy.

 The system is the part of the world in which we have a special interest. The surroundings
is the region outside the system where we make our measurements.
-An open system has a boundary through which matter can be transferred. A
closed system has a boundary through whichmatter cannot be transferred. An isolated
system has a boundary through which neither matter nor energy can be transferred.
 Energy is the capacity to do work. The internal energy is the total energy of a system.
 Work is the transfer of energy by motion against an opposingrce, dw = −Fdz . Heat is the
transfer of energy as a result of a temperature difference between the system and the
surroundings.
 An exothermic process releases energy as heat to the surroundings.
 An endothermic process absorbs energy as heat from the surroundings.
 A state function is a property that depends only on the current state of the system and is
independent of how that state has been prepared.
 The First Law of thermodynamics states that the internal energy of an isolated system is
constant, ∆U = q + w.
 Internal energy ,heat,and work are all measured in the same units, the Joule (J)
 . Expansion work is the work of expansion (or compression) of a system, dw = −PexdV.
 The work of free expansion is w = 0.
 The work of expansion against a constant external pressure is w = −Pex∆V.
Vf
 The work of isothermal reversible expansion of aperfect gas is w = −nRT ln( )
Vi
 A reversible change is a change that can be reversed by an infinitesimal modification of
a variable
 Maximum work is achieved in a reversible change.
 Calorimetry is the study of heat transfers during physical and chemical processes.
∂U
 The heat capacity at constant volume is defined as CV = ( ).
∂T V
∂H
 The heat capacity at constant pressure is Cp= ( )p.
∂T
 For a perfect gas, the heat capacities are related by Cp − CV = nR.
 The enthalpy is defined as H = U + pV.
 The enthalpy change is the energy transferred as heat at constant pressure, ∆H = qp.
 During a reversible adiabatic change, the temperature of a perfect gas varies according to
Vi
Tf = Ti( )1/c, c = CV,m/R.
Vf
Cp. m
 The pressure and volume are related by pV γ = constant, with γ =( ).
Cv . m
 The standard enthalpy change is the change in enthalpy for a process in which the initial
and final substances are in their standard states.
 The standard state is the pure substance at1 bar
 Enthalpy changes are additive, as in ∆subHø= ∆fusHø + ∆vapHø.
 The enthalpy change for a process and its reverse are related by ∆forwardHø = −∆reverseHø.
 Hess’s law states is the sum of the standard enthalpies of the individual reactions into
which a reaction may be divided.
 The standard enthalpy of formation (∆fHø) is the standard reaction enthalpy for the
formation of the compound from its elements in their reference states. The reference state
is the most stable state of an element at the specified temperature and 1 bar.
 The standard reaction enthalpy may be estimated by combining enthalpies of formation,
∆rH 7 ø= ∑Productsν∆fHø− ∑Reactantsν∆fHø.
 The temperature dependence of the reaction enthalpy is given by Kirchhoff’s law,
T2
ø ø
∆rH (T2) = ∆rH (T1) +∫ ∆ rC pdT.
T1
 An exact differential is an infinitesimal quantity that, when integrated, gives a result that
is independent of the path between the initial and final states. An inexact differential
is an infinitesimal quantity that, when integrated, gives a result that depends on the path
between the initial and final states.
∂U
 The internal pressure is defined as πT = ( )T . For aperfect gas, πT = 0.
∂V
 The Joule–Thomson effect is the cooling of a gas by isenthalpic expansion.
∂T
 The Joule–Thomson coefficient is defined as µ = ( )H.
∂P
∂T
 The isothermal Joule–Thomson coefficient is defined as µT = ( ) = −Cpµ.
∂P T
 The inversion temperature is the temperature at which the Joule–Thomson coefficient
changes sign.

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THE FIRST LAW (DISCUSSION)

 Thermodynamics is the study of the transformations of energy.

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