ThermodynamicsScience of energy deals with conversion of energy fromone form to another.EnergyThe amount of work that can be performed by a force toproduce change.The law of conservation of energy.Energy can neither be created (produced) nor destroyed byitself. It can only be transformed from one form to another.According to conservation law the total inflow of energyinto a system must equal the total outflow of energy fromthe system, plus the change in the energy contained withinthe system
Thermodynamic system.In thermodynamics, a thermodynamic system, originallycalled a working substance, is defined as that part of theuniverse that is under consideration. Anything underconsideration is called a system. A boundary separatesthe system from the rest of universe, which referred toas environment, surroundings, or reservoir. A usefulclassification of thermodynamics systems is based onthe nature of the boundary and the quantities flowingthrough it, such as matter, energy, work, heat.
Isolated systemIn isolate system, any modification of the environment has noeffect on the system, and any change of the system has no effecton the environment neither matter nor energy in any form areexchanged between the system. An example of an isolatedsystem would be an insulated rigid container, such as aninsulated gas cylinder.Closed system.A closed system is a system which may exchange energy in anyform with environment ( work, heat,….) but which cannotexchange any matter.Open system.An open system can exchange both energy and matter with itsenvironment. The ocean would be an example of an opensystem.
Zeroth law of thermodynamics, about thermalequilibrium.If two thermodynamics system are separately inthermal equilibrium with a third, they are also inthermal equilibrium with each other.Properties of system.1- Pressure (p).2- Temperature (T) or (t).3- Volume (V).4- Mass (m).Some are defined in term of other oneDensity : mass per unit volume. ρ=m/V = (Kg/ m3)Relative density ρs = ρ/ρH2o dimension less quantity.Specific volume: volume per unit mass .ν= V/m = 1/ρ= m3/kg.
Thermodynamics coordinate.(P, V, and T) or (P, ρ, and T ) are thermodynamiccoordinate any change of these three coordinate gavethermodynamic process.1- An isobaric process occurs at constant pressure.2- An isochoric process occurs at constant volume.3- An isothermal process occurs at constanttemperature.
EquilibriumA state of balance. A system which is inequilibrium expressing no change from it’ssurrounding. There are many types of equilibrium•Two systems are in thermal equilibrium whentheir temperatures are the same (no temp.difference).• Two system are in mechanical equilibrium whentheir pressures are the same (no pressure changewith time).• Two system in chemical equilibrium when theirchemical potentials are the same (no chemicalreaction occurs).
Adiabatic process. In thermodynamics, an adiabatic process is a thermodynamicprocess in which no heat is transferred to or from the system.(Thequantity of thermal energy is constant). Q1=Q2 dQ=0 Q= constant. Temperature scales. C0= (F0 - 32)*5/9 C0= K - 273.15 F0 = R0 – 459.67 Note ΔK0=ΔC0 ΔR0 = ΔF0 C0 = degree Celsius. F0 = degrees Fahrenheit. R0 = Rankine . K= Kelvin.
Equation of statean equation of state is a relation between statevariables describing the state of matter undera given set of physical conditions. It is aconstitutive equation which providesmathematical relationship between two ormore state functions associated with thematter, such as its (T, P and V).
Boyle’s law. When gas is kept at constant temperature itspressure is inversely proportion to the volume. P1 V1 = P2 V2 Charles’s law. When the pressure of the gas kept constant thevolume directly proportional to the temperature. V1/T1 = V2/T2 The General Law P1V1/T1 = P2V2/T2 Combined Boyle’s law and Charles’s law into thefirst statement of ideal gas law. PV=nRT n=number of moles R= gas constant= 8.314 (j/mole.K) T = temp. In (K) .
Real gas equation of state Real gas has distance between moleculesless than distance between ideal gas molecule. In real gas we have two equation. 1- Glasius equation. P(V- nb)=nRT 2- Vander Waals equation. (P + a/V3)(V- b)= RT Where a and b are constant for any one gasbut differ for different gasses.
HeatThe form of energy that is transferred betweentwo systems by temperature difference. Heattransferred during the process between two states(state1 and state 2) is denote by Q12 or Q. Heathas energy unit KJ or J.q=Q/m (KJ/Kg). heat transfer per unit mass.
WORK .Work W is performed whenever a force acts through a distance.By definition, the quantity of work is given by the equation:Where F is the component of force acting along the line of thedisplacement dl. When integrated, this equation yields the workof a finite process. By convention, work is regarded as positivewhen the displacement is in the same direction as the appliedforce and negative when they are in opposite directions.
Figure (**) shows apath for compression of a gas from point 1with initial volume V1t at pressure P1 to point 2 with volumeVt2 at pressure P2. This path relates the pressure at any pointof the process to the volume. The work required is given byEq. (*) and is proportional to the area under the curve of Fig.(**). The unit of work is the newton-meter or joule, symbol J.
THE FIRST LAW OF THERMODYNAMICSThe recognition of heat and internal energy as forms of energymakes possible a generalization of the law of conservation ofmechanical energy include heat and internal energy in addition towork and external potential and kinetic energy. Indeed, thegeneralization can be extended to still other forms, such as surfaceenergy, electrical energy, and magnetic energy. This generalizationwas at first a postulate. However, the overwhelming evidenceaccumulated over time has elevated it to the stature of a law ofnature, known as the first law of thermodynamics. One formalstatement is:Although energy assumes many forms, the total quantity ofenergy is constant, and when energy disappears in one form itappears simultaneously in other forms.
In application of the first law to a given process, the sphere ofinfluence of the process is divided into two parts, the system andits surroundings. The region in which the process occurs is setapart as the system; everything with which the system interactsis the surroundings. The system may be of any size dependingon the application, and its boundaries may be real or imaginary,rigid or flexible. Frequently a system consists of a singlesubstance; in other cases it may be complex. In any event, theequations of thermodynamics are written with reference to somewell-defined system. This focuses attention on the particularprocess of interest and on the equipment and material directlyinvolved in the process. However, the first law applies to thesystem and surroundings, and not to the system alone. In itsmost basic form, the first law requires:
Where the difference operator "Δ" signifies finite changes in thequantities enclosed in parentheses. The system may change in itsinternal energy, in its potential or kinetic energy, and in thepotential or kinetic energy of its finite parts. Since attention isfocused on the system, the nature of energy changes in thesurroundings is not of interest.In the thermodynamic sense, heat and work refer to energy intransit across the boundary which divides the system from itssurroundings. These forms of energy are not stored, and are nevercontained in a body or system. Energy is stored in its potential,kinetic, and internal forms; these reside with material objects andexist because of the position, configuration, and motion of matter.
The choice of signs used with Q and W depends on whichdirection of transport is regarded as positive. Heat Q and work Walways refer to the system, and the modern sign convention makesthe numerical values of both quantities positive for transfer intothe system from the surroundings. The corresponding quantitiestaken with reference to the surroundings, Qsurr and Wsurr havtheopposite sign, i.e., Qsurr = - Q and Wsurr = - W. With thisunderstanding:Closed systems often undergo processes that cause no change in the systemother than in its internal energy. For such processes, reduces to:
HEATCAPACITY We remarked earlier that heat is often viewed in relation to its effect on the object to which or from which it is transferred. This is the origin of the idea that a body has a capacity for heat. The smaller the temperature change in a body caused by the transfer of a given quantity of heat, the greater its capacity. Indeed, a heat capacity might be defined:The difficulty with this is that it makes C, like Q, a process-dependentquantity rather than a state function. However, it does suggest the possibilitythat more than one useful heat capacity might be defined. In fact two heatcapacities are in common use for homogeneous fluids; although their namesbelie the fact, both are state functions, defined unambiguously in relation toother state functions.
Heat Capacity at Constant VolumeHeat Capacity at Constant Pressure Isothermal Process Isobaric Process