University Physics Volume 2

# Summary

### 3.1Thermodynamic Systems

• A thermodynamic system, its boundary, and its surroundings must be defined with all the roles of the components fully explained before we can analyze a situation.
• Thermal equilibrium is reached with two objects if a third object is in thermal equilibrium with the other two separately.
• A general equation of state for a closed system has the form $f(p,V,T)=0,f(p,V,T)=0,$ with an ideal gas as an illustrative example.

### 3.2Work, Heat, and Internal Energy

• Positive (negative) work is done by a thermodynamic system when it expands (contracts) under an external pressure.
• Heat is the energy transferred between two objects (or two parts of a system) because of a temperature difference.
• Internal energy of a thermodynamic system is its total mechanical energy.

### 3.3First Law of Thermodynamics

• The internal energy of a thermodynamic system is a function of state and thus is unique for every equilibrium state of the system.
• The increase in the internal energy of the thermodynamic system is given by the heat added to the system less the work done by the system in any thermodynamics process.

### 3.4Thermodynamic Processes

• The thermal behavior of a system is described in terms of thermodynamic variables. For an ideal gas, these variables are pressure, volume, temperature, and number of molecules or moles of the gas.
• For systems in thermodynamic equilibrium, the thermodynamic variables are related by an equation of state.
• A heat reservoir is so large that when it exchanges heat with other systems, its temperature does not change.
• A quasi-static process takes place so slowly that the system involved is always in thermodynamic equilibrium.
• A reversible process is one that can be made to retrace its path and both the temperature and pressure are uniform throughout the system.
• There are several types of thermodynamic processes, including (a) isothermal, where the system’s temperature is constant; (b) adiabatic, where no heat is exchanged by the system; (c) isobaric, where the system’s pressure is constant; and (d) isochoric, where the system’s volume is constant.
• As a consequence of the first law of thermodymanics, here is a summary of the thermodymaic processes: (a) isothermal: $ΔEint=0,Q=W;ΔEint=0,Q=W;$ (b) adiabatic: $Q=0,ΔEint=−W;Q=0,ΔEint=−W;$ (c) isobaric: $ΔEint=Q−W;ΔEint=Q−W;$ and (d) isochoric: $W=0,ΔEint=Q.W=0,ΔEint=Q.$

### 3.5Heat Capacities of an Ideal Gas

• For an ideal gas, the molar capacity at constant pressure $CpCp$ is given by $Cp=CV+R=dR/2+RCp=CV+R=dR/2+R$, where d is the number of degrees of freedom of each molecule/entity in the system.
• A real gas has a specific heat close to but a little bit higher than that of the corresponding ideal gas with $Cp≃CV+R.Cp≃CV+R.$

### 3.6Adiabatic Processes for an Ideal Gas

• A quasi-static adiabatic expansion of an ideal gas produces a steeper pV curve than that of the corresponding isotherm.
• A realistic expansion can be adiabatic but rarely quasi-static. Do you know how you learn best?
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