ThermodynamicsScience of energy deals with conversion of energy from one form to another. Energy The amount of work that can be performed by a force to produce change. The law of conservation of energy. Energy can neither be created (produced) nor destroyed by itself. It can only be transformed from one form to another. According to conservation law the total inflow of energy into a system must equal the total outflow of energy from the system, plus the change in the energy contained within the system

Thermodynamic system. In thermodynamics, a thermodynamic system, originally called a working substance, is defined as that part of the universe that is under consideration. Anything under consideration is called a system. A boundary separates the system from the rest of universe, which referred to as environment, surroundings, or reservoir. A useful classification of thermodynamics systems is based on the nature of the boundary and the quantities flowing through it, such as matter, energy, work, heat.

Surroundings

Isolated system In isolate system, any modification of the environment has no effect on the system, and any change of the system has no effect on the environment neither matter nor energy in any form are exchanged between the system. An example of an isolated system would be an insulated rigid container, such as an insulated gas cylinder. Closed system. A closed system is a system which may exchange energy in any form with environment ( work, heat,….) but which cannot exchange any matter. Open system. An open system can exchange both energy and matter with its environment. The ocean would be an example of an open system.

Zeroth law of thermodynamics, about thermal equilibrium. If two thermodynamics system are separately in thermal equilibrium with a third, they are also in thermal 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 one Density : 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 thermodynamic coordinate any change of these three coordinate gave thermodynamic process. 1- An isobaric process occurs at constant pressure. 2- An isochoric process occurs at constant volume. 3- An isothermal process occurs at constant temperature.

Equilibrium A state of balance. A system which is in equilibrium expressing no change from it’s surrounding. There are many types of equilibrium•Two systems are in thermal equilibrium when their temperatures are the same (no temp. difference). • Two system are in mechanical equilibrium when their pressures are the same (no pressure change with time). • Two system in chemical equilibrium when their chemical potentials are the same (no chemical reaction occurs).

Adiabatic process. In thermodynamics, an adiabatic process is a thermodynamic

process in which no heat is transferred to or from the system.(The quantity of thermal energy is constant).

Q1=Q2dQ=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 state an equation of state is a relation between state variables describing the state of matter under a given set of physical conditions. It is a constitutive equation which provides mathematical relationship between two or more state functions associated with the matter, such as its (T, P and V).

Boyle’s law. When gas is kept at constant temperature its

pressure is inversely proportion to the volume. P1 V1 = P2 V2 Charles’s law. When the pressure of the gas kept constant the

volume 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 the

first 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 molecules

less 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 gas

but differ for different gasses.

Heat The form of energy that is transferred between two systems by temperature difference. Heat transferred during the process between two states (state1 and state 2) is denote by Q12 or Q. Heat has 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 the displacement dl. When integrated, this equation yields the work of a finite process. By convention, work is regarded as positive when the displacement is in the same direction as the applied force and negative when they are in opposite directions.

or, since A is constant,

Integrating,

Figure (**) shows apath for compression of a gas from point 1 with initial volume V1t at pressure P1 to point 2 with volume Vt2 at pressure P2. This path relates the pressure at any point of the process to the volume. The work required is given by Eq. (*) 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 THERMODYNAMICS The recognition of heat and internal energy as forms of energy makes possible a generalization of the law of conservation of mechanical energy include heat and internal energy in addition to work and external potential and kinetic energy. Indeed, the generalization can be extended to still other forms, such as surface energy, electrical energy, and magnetic energy. This generalization was at first a postulate. However, the overwhelming evidence accumulated over time has elevated it to the stature of a law of nature, known as the first law of thermodynamics. One formal statement is: Although energy assumes many forms, the total quantity of energy is constant, and when energy disappears in one form it appears simultaneously in other forms.

In application of the first law to a given process, the sphere of influence of the process is divided into two parts, the system and its surroundings. The region in which the process occurs is set apart as the system; everything with which the system interacts is the surroundings. The system may be of any size depending on the application, and its boundaries may be real or imaginary, rigid or flexible. Frequently a system consists of a single substance; in other cases it may be complex. In any event, the equations of thermodynamics are written with reference to some well-defined system. This focuses attention on the particular process of interest and on the equipment and material directly involved in the process. However, the first law applies to the system and surroundings, and not to the system alone. In its most basic form, the first law requires:

Where the difference operator "Δ" signifies finite changes in the quantities enclosed in parentheses. The system may change in its internal energy, in its potential or kinetic energy, and in the potential or kinetic energy of its finite parts. Since attention is focused on the system, the nature of energy changes in the surroundings is not of interest. In the thermodynamic sense, heat and work refer to energy in transit across the boundary which divides the system from its surroundings. These forms of energy are not stored, and are never contained in a body or system. Energy is stored in its potential, kinetic, and internal forms; these reside with material objects and exist because of the position, configuration, and motion of matter.

The choice of signs used with Q and W depends on which direction of transport is regarded as positive. Heat Q and work W always refer to the system, and the modern sign convention makes the numerical values of both quantities positive for transfer into the system from the surroundings. The corresponding quantities taken with reference to the surroundings, Qsurr and Wsurr havthe opposite sign, i.e., Qsurr = - Q and Wsurr = - W. With this understanding:

Closed systems often undergo processes that cause no change in the system other 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-dependent quantity rather than a state function. However, it does suggest the possibility that more than one useful heat capacity might be defined. In fact two heat capacities are in common use for homogeneous fluids; although their names belie the fact, both are state functions, defined unambiguously in relation to other state functions.

Heat Capacity at Constant Volume

Heat Capacity at Constant Pressure

Isothermal Process

Isobaric Process

lsochoric (Constant- V) Process