Some thermodynamical terms

Some common terms, which are frequently used in the discussion of energetics need to be known.


'A specified part of the universe which is under investigation is called the system'. The system is separated from the rest of the universe by a definite (real or imaginary) boundary.

Types of Systems

Systems are of various types depending upon the exchange of mass and energy and the constituent between the system and the surroundings.

Isolated system

A system, which can neither exchange mass nor energy with surroundings, is called an isolated system. They are also called insulated systems. Hot coffee (in contact with its vapour) in a closed and insulated thermos is an example of isolated system. Since the vessel is closed, matter can neither enter nor leave the vessel. Moreover, as the vessel is well insulated, heat can neither leave the system nor enter from the surroundings.

Closed system

A system, which can exchange energy but not mass with the surroundings is called a closed system. For example, boiling water in a closed steel vessel is an example of a closed system. The energy can be gained or lost (through the steel walls) but not matter. Similarly, all reactions carried out in a closed container are examples of closed systems.

Open system

A system, which can exchange matter as well as energy with the surroundings is called an open system. All reactions carried out in open containers are examples of open systems. Evaporation of water in a beaker or hot coffee in a cup represents an open system. Here vapour of water or coffee (matter) can leave the system and escape into atmosphere. The heat energy required for this purpose is absorbed from the surroundings. All physical and chemical processes taking place in open in our daily life are open systems because they are continuously exchanging matter and energy with the surroundings.

Homogeneous system

A system is called homogeneous if physical properties and chemical composition are identical throughout the system. A pure gas or consistent mixture of gases e.g., an oxygen cylinder, or a pure liquid or solid in a container are examples of homogeneous systems.

Heterogeneous system

A system is said to be heterogeneous if it consists of parts separated by definite boundaries, each of which has different physical and chemical properties. A mixture of fertilizer granules (N P K) or ice with water are typical examples of this system.

State of a System

The condition of the system when its macroscopic properties have definite measurable values, is the state of a system. If any of the macroscopic properties of the system changes such as its temperature, pressure etc., the state of the system is also said to change. To define the state of the system all its macroscopic properties need not be specified.

State variable

A measurable property required to describe the state of a system is called state variable. For example, temperature (T), pressure (P), volume (V) etc. are the state variables. For any system, a certain minimum number of variables are sufficient to define its state, as the other variables become automatically fixed.

For instance, in a system consisting of an ideal gas, the state may be defined by only three variables such as temperature (T), pressure (P) and volume (V). The values of other variables such as amount of gas, density, etc. will be definite and can be easily calculated.

State functions

A property whose values depend on only the state of the system and not on the path by which the change from initial to final state is brought about is called state function. The change in the value of the state functions depends only upon the initial and final state the system.

Some common state functions used in thermodynamics are, pressure, (P), volume (V), temperature (T), internal energy (E or U), enthalpy (H), entropy (S) free energy (G) etc.

State properties

State variables and functions may posses either extensive or intensive properties.

Extensive properties are those, whose values depend on the quantity of matter contained in any system or size of the system. Typical properties are mass, volume, area, energy, number of moles, enthalpy, entropy etc.

Intensive properties are those, which depend on the nature of the substance system but is independent of its amount / size in the system. Examples are density, viscosity, surface tension, temperature, pressure, boiling point etc.

State processes

Isothermal process is that in which the temperature of the system remains constant.

Adiabatic process is where there is no exchange of heat between the system and the surrounding.

Isobaric process is when the pressure on the system remains constant during any operation.

Isochroic process is that in which the volume of the system does not change.

Cyclic process is one when the system returns to its initial state after having undergone a series of change.

Reversible process occurs when a process is carried out slowly so that the system and the surrounding are always in equilibrium.

Irreversible process is that which is carried out rapidly so that the system does not return to its initial state.


The region outside the boundary of the system is termed as surrounding. The rest of the universe, which is not the part of the system, is separated by a real and imaginary boundary. The boundary sets the limits of the system.

For example, while studying the effects of adding acid to water in a beaker, the acid water mixture forms the system, the walls of the container is the boundary and the rest of the part that is around it is the surrounding.

Energy Terms in Thermodynamics

Every system has a definite amount of energy. It can exchange energy (lose or gain) with the surroundings in a variety of ways. In chemical systems the two important modes of transference of energy between the system and the surroundings are heat and work.

Heat (Q)

Energy exchanged between the system and the surroundings when they are at different temperatures is commonly known as heat. If a system is at a higher temperature than the surroundings, then there is a flow of heat (or energy) from system to surroundings, causing a decline of the systems temperature and an increase in the surrounding temperature. These processes continue till the fall in temperature of the system and rise in temperature of the surroundings, become equal.

Heat exchange between system and surroundings

Fig: 5.1 - Heat exchange between system and surroundings

Work (W)

Work is another mode of transference of energy. Work is said to be done if the point of application of force undergoes displacement in the direction of the force. If the system loses energy, we say that the work is done by the system. Alternatively, if the system gains energy, we say that work is done on the system.

For example, if a gas, enclosed in a cylinder with a piston, has a higher pressure than the surroundings, the piston will move upward until the pressure inside and outside become equal. The gas expands against a constant external pressure 'P' and its volume changes by an amount equal toDV. The energy transfer that takes in this case is called work. At this step the work is done by the system on the surroundings. This is given as:

Work done by the system = PDV.

Alternatively, if the system is at lower pressure, piston will be pushed down until the pressure of the system becomes equal to that of the surroundings. In this case work is done on the system by the surroundings.

Obviously, if there is no change in volume, i.e., DV = 0, no work is done by the system, i.e. work = 0.

In addition to these two modes, radiant energy and electrical energy are also modes of transference of energy between the system and the surroundings.

Units of heat and work

The heat changes are measured in calories (cal), kilo calories (kcal), joules (J) or kilojoules (kJ). These are related as:

1 cal = 4.184J

1 kcal = 4.184kJ

The S.I. unit of heat is joule or kilojoule

Work is measured in terms of ergs or joules.

I Joule = 107 ergs.

1calorie = 4.184 x 107 ergs

The S.I. unit of work is joule.

Sign Conventions for Heat and Work

The signs of 'w' and 'q' are related to the internal energy change. When 'w' or 'q' is positive, it means that energy has been supplied to the system as work or as heat. The internal energy of the system in such a case increases. On the other hand, if 'w' or 'q' is negative, it means that energy has left the system as work or heat. The internal energy of the system decreases. The signs of 'q' and 'w' are:

Heat absorbed by the system= q positive

Heat evolved by the system= q negative

Work done on the system= w positive

Work done by the system= w negative


1. Assuming ideal behaviour, calculate the work done when 1.6 mole of water evaporates at 373 K against the atmospheric pressure of 760 mm of Hg.


Volume of 1.6 mole of water at 373 K in gaseous state

Volume of 1 mol = 18 g of liquid water (density = 1 g ml-1)

= 18 x 1.6 x 10-3L = 0.0288 L

Now work done (W) = -P(V2 - V1)

= -1(48.93 - 0.0288) = - 48.90 atm L

= - 48.90 x 101.325 J = - 4954.8 J.

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