The state of a thermodynamic system can be changed by a process. In other wards, a process gives the path or operation by which a system changes from one state to another. The process may be accompanied by an exchange of matter and energy between the system and surroundings.
Type of thermodynamic Process:
Some common types of thermodynamic process are as follows-
- Isothermal process
- Adiabatic process
- Isobaric process
- Isochoric process
- Reversible process
- Irreversible process
- Cyclic process
1) Isothermal process:-A process in which the temperature of system remains constant is called isothermal process. In such a system, heat is either supplied to the system or removed from it.
2) Adiabatic process: – A process in which the system does not exchange heat with the surroundings i.e. no heat leaves or enters the system. In such a process, temperature of the system always changes. The systems in which such processes occur are thermally insulated from the surroundings.
3) Isobaric process: – A process in which the pressure of the system remain constant.
4) Isochoric process: – A process in which the pressure of the system remain constant.
5) Reversible process: – A process in which the direction may be reversed at any stage by merely a small change in a variable like temperature, pressure, etc.
6) Irreversible process: – A process which is not reversible is called irreversible process. Irreversible process proceeds only one direction. e.g. all natural process is irreversible process.
7) Cyclic process: – A process in which the system undergoes a series of changes and ultimately returns to its original state is called a cyclic process.
Internal Energy (E):
Every substance possesses a definite quantity of energy which depends upon factors such as chemical nature of the substance, temperature and pressure. This is known as intrinsic energy or internal energy and is represented by the symbol E. It is made up of kinetic energy and potential energy of the constituent particles. The kinetic energy arises due to motion of its particles and includes their translational energy, rotation energy, vibrational energy etc. The potential energy arises from different types of interactions between the particles and includes electronic energy, energy due to molecular interactions, nuclear energy, etc. The sum of all forms of energies stored in atoms or molecules is internal energy.
Different substances have different internal energies depending upon the nature of the constituting atoms, bonds and other conditions of temperature, pressure, etc. For example, the internal energy of 1 mole of carbon dioxide will be different from the internal energy of 1 mole of sulphur dioxide even under similar conditions of temperature and pressure. Further, the internal energy of 1 mole of water at 300 K is different than that of one mole of water at 310 K under same atmospheric pressure.
Internal Energy Change (∆E):
The absolute value of internal energy possessed by a substance cannot be calculated because it is not possible to find, the accurate values of different types of energies, such as translational, vibrational, rotational, chemical energy, etc. stored in a system. However, we are interested mostly in change in internal energy which occurs during chemical reactions. The change in internal energy of a reaction may be considered as the difference between the internal energies of the two states.
Let EA and EB are the internal energies in states A and B respectively. Then the difference between the internal energies in the two states will be
E = EB — EA
The difference in internal energies (AU) has a fixed value and will be independent of the path taken between two states A and B. For chemical reactions, the change in internal energy may be considered as the difference between the internal energies of the products and that of the reactants, i.e.
∆E= E products — E reactants
= Ep – Er
Where Ep is the internal energy of the products, Er is internal energy of the reactants and ∆E gives the change in internal energy.
If the internal energy of the products is less than the internal energy of the reactants, then ∆E be negative.
∆E=Ep -Er = -ve (Ep< Er)
On the other hand, if the internal energy of the products is more than the internal energy reactants, the ∆E would be positive.
∆E=Ep –Er= +ve (Ep>Er)
Thus, the internal energy, E is a state function. This means that ∆E depends only on the initial and final states and is independent of the path. In other words, ∆E will be same even if the change is brought about differently.
Heat (Q):
Energy is exchanged between the system and the surroundings as heat if they are at different temperatures. e.g. if a system is at a higher temperature than the surroundings, then energy (or heat) is lost to the surroundings causing a fall in temperature of the system and rise in temperature of the surroundings. This process continues till the temperatures of the system and the surroundings become equal. If the temperature of the system is lower than that of the surroundings, then energy is gained by the system. In words, we say that there is a flow of heat from system to surroundings if system has higher temperature and there is flow of heat from surroundings to the system if the former is at higher temperature.
This exchange of energy which is a result of temperature difference is called heat (q). The q is positive when heat is transferred from the surroundings to the system and q will be negative when heat is transferred from system to surroundings.
Work (w):
It is another mode of transference of energy. Work is said to be performed if the point of application of force is displaced in the direction of the force. It is equal to the force multiplied by the displacement (distance through which the force acts).
There are two main types of work which we generally come across. These are (i) electrical work and (ii) mechanical work. Electrical work is important in systems where reaction takes place between ions whereas mechanical work is performed when a system changes its volume in the presence of external pressure. Mechanical work is important specially in systems that contain gases. This is also known as pressure-volume work.
Note:
(i) If the gas expands, Vf >Vi then work is done by the system and w is negative.
(ii) If the gas contracts, Vf <Vi then work is done on the system and w is positive.
Heat Capacity:
The heat capacity of a sample of substance is the quantity of heat energy required to raise its temperature by 1K (or 1◦C). It can be represented as-
C = Q/ ∆t
Where, Q = the quantity of heat given to the sample.
∆t = the raise in the temperature of the substance.
Specific Heat Capacity:
Specific heat capacity is also called specific heat. The specific heat capacity of a substance is equal to the quantity of heat required to raise the temperature of 1 gram of substance by 1◦C. The specific heat capacity of a substance is equal to the ratio of heat capacity of the substance to mass of the substance.
Molar Heat Capacity:
The molar heat capacity (m) of a sample of substance is defined as the quantity of heat energy required to raise the temperature of one mole of a substance by 1K (or 1◦C). Thus, Molar heat capacity = specific heat capacity x molar mass.
Enthalpy (H):
Chemical reactions are generally carried out at constant pressure i.e. at atmospheric pressure. ∆E gives the change in internal energy at constant volume. To express the energy change at constant pressure, a new terms called enthalpy change (∆H) was used. Enthalpy may be defined as the sum of the internal energy and the products of its pressure and volume of the system. Enthalpy is denoted by the letter H and is given as-
H = E + PV
Since, E and PV are state functions; hence H also is a state function. Every substance has a certain amount of enthalpy. Enthalpy of any substance depends upon its temperature and pressure.
Enthalpy Change (∆H):
Every substance has a definite value of enthalpy in a particular state. Like internal energy, the absolute value of enthalpy cannot be measured. However, the change in enthalpy accompanying a process can be determined. Thus for a process represented as:
A ——–> B
The change in enthalpy may be expressed as:
∆H = HB – HA
Where, HA is the enthalpy of state A and HB is the enthalpy of state B and ∆H is the enthalpy change. For a chemical reaction, ∆H can be given as difference between the enthalpies of the products and the reactants i.e.
∆H = H products – H reactants
∆H = Hp– Hr
Where Hp is the enthalpy of the products, Hr is the enthalpy of the reactants and ∆H is the enthalpy change.
If a reaction is carried out at constant temperature and constant pressure, the heat exchanged by the system with the surroundings (i.e., heat change) is equal to change in enthalpy. Thus, the enthalpy change of a reaction is equal to the heat absorbed or evolved during a reaction at constant temperature and pressure.
∆H= Heat evolved or absorbed in a reaction at constant temperature and constant pressure
It may be noted that the amount of heat exchanged with the surroundings for a reaction at constant pressure (∆H) is different from that exchanged at constant volume (∆E) and temperature. This can be easily understood. At constant pressure, the volume of the reacting system changes. If the volume increases, the system expands against the atmospheric pressure and energy is required for this expansion. Therefore, a part of energy will be used for the expansion. As a result, the amount of heat exchanged at constant pressure (∆H) would be less than the amount of heat exchanged at constant volume (∆E).
Alternatively, if the system contracts at constant pressure, work is done on the system and the system absorbs some energy from the surroundings. Therefore, the amount of heat exchanged at constant pressure is greater than that exchanged at constant volume.