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System, Surroundings and Some

In this section we shall study some basic terms/concepts which are needed in the study of chemical energetics.

SYSTEM

A system is defined as that part of the Universe which is under investigation. For example, if we are interested in investigating the neutralization of NaOH with HCl, the solutions of sodium hydroxide and hydrochloric acid constitute the system. Similarly, if we are studying the effect of temperature on the properties of water, then water will be taken as the system.

SURROUNDINGS

The part of the Universe other than the system is known as surroundings.

Thus, Universe = System + Surroundings

The entire universe, other than the system is not affected by the changes taking place in the system. Therefore, for all practical purposes, the surroundings are that portion of the remaining universe which can interact with the system. Generally, the region of space in the neighbourhood of the system constitutes its surroundings.

In order to keep track of exchange of matter and energy between system and surroundings, it is necessary to think of the system as separated from the surroundings by some sort of wall which may be real or imaginary. The wall which separates the system from the surroundings is called boundary. For example, if a reaction mixture is taken in a beaker, the reaction mixture constitutes the system, the walls of the beaker constitute the boundary and everything else is the surroundings.

 

TYPES OF SYSTEM

Systems may be classified into three types on the basis of movement of matter and energy in and out of the system:

(a) Open system          (b) Closed system

(c) Isolated system.

 (a) Open System. A system which can exchange mass as well as energy with the surroundings, is called an open system. For example, heating of calcium carbonate in an open vessel. In this case heat is supplied to the system by the burner, while CO2 escapes into the surroundings.

(b) Closed System. A system which can exchange energy with the surroundings but not mass is called a closed system. For example, calcination of CaCO3 taken in sealed bulb. On heating, CaCO3 decomposes into CaO and CO2, however, CO2 cannot escape and remains trapped in the bulb.

(c) Isolated System. A system which can neither exchange mass nor energy with the surroundings is called an isolated system. For example, a reaction carried out in a closed, well insulated container (such as thermous flask). The open, closed and isolated systems have been shown in Fig. 16.1.

MACROSCOPIC SYSTEM

 

A system containing a large number of chemical species (atoms, ions or molecules) is called macroscopic system. In fact, the word macroscopic conveys the sense of appreciable

quantities. The properties of the system which arise from the collective behaviour of large number of species are called macroscopic properties. These properties are pressure temperature, volume, composition, refractive index, surface tension, viscosity, density, etc. For example, in order to note the temperature of water we do not deal with individual molecules but we consider the molecules in bulk.

 

EXTENSIVE AND INTENSIVE PROPERTIES

The various physical properties of the system may be classified into two types:

 

(a) Extensive Properties. The properties of the system which depend upon the quantity or size of matter present in it are called extensive properties. Some common examples of these properties are mass, volume, internal energy, enthalpy heat capacity, etc.

 

(b) Intensive Properties. The properties of the system which are independent of the quantity or size of matter present in it are called intensive properties. Some common examples are temperature, pressure, refractive index, viscosity, specific heat, density, etc.

 

It may be noted that the ratio of the two extensive properties become intensive in nature. For example, mass and volume are extensive properties, but the ratio mass/volume, i.e., density is independent of quantity of matter and is intensive. Similarly, heat capacity is extensive, but molar heat capacity is intensive. In general, if X is any extensive property of n mol of system, then the molar property of the system, Xm is intensive because it refers to the property of 1 mol of the system and is independent of the quantity of matter. Xm = x/n.

 

STATE OF THE SYSTEM

State of the system refers to the conditions of existence of a system when its macroscopic properties have definite values. In order to make useful calculations, it is necessary to describe the system completely before and after it undergoes any change the state of the system can be described by specifying a certain minimum number of properties such as pressure (P), volume (V), temperature (T) and amount (n) or composition. Once these minimum number of macroscopic properties are fixed, the other properties automatically acquire definite values. Some notable features about the thermodynamic state are:

 

  • Variation in one or more macroscopic properties brings a change in the state of the system, when other macroscopic properties attain new values. The macroscopic properties are thus, called state variables or state functions.
  • Initial state refers to the starting state of system in equilibrium. After interaction with surroundings (involving exchange of matter or energy or both) the system attains another equilibrium state which is referred to as a final state of the system.

 

  • Thermodynamic state of the system must not be confused with physical state or phase.

 

  • A system is said to be in thermodynamic equilibrium state if its macroscopic properties do not change with time.

 

STATE FUNCTIONS

The properties whose values depend only upon the initial. and final states of the system and are independent of the manner as to how the change is brought about, are called state functions. The concept of state function can be easily understood from the following analogy. If we consider ‘h’ as the height between the top and bottom of the mountain, then ‘h’ is independent of the path followed in reaching the top of the mountain. Here, the parameter h is analogous to state function. In thermodynamics, some common state functions are internal energy (U), enthalpy (H), entropy (S), Gibb’s energy (G), pressure (P), temperature (T), volume (V), etc.

It may be noted that two very important thermodynamic parameters namely; heat (q) and work (w) are not the state functions because they are path dependent.