Thermodynamics Common Terms, State Function, Reversible & Irreversible

Thermodynamics is about understanding how heat and energy work. It looks at how different types of energy can change from one form to another. There are three main rules that explain how energy behaves. The term “thermodynamics” was coined by William Thomson (Lord Kelvin) in 1849. Let us now explore some basic terms you will come across in thermodynamics.

What is Thermodynamics?

The branch of chemistry that deals with the movement of energy from one form to the other and the relation between heat and temperature with energy and work done is called thermodynamics.

• Molecules store chemical energy, which can be liberated as heat when fuels such as methane, LPG, or coal burn in the presence of air. This released energy can be used to perform mechanical work in engines or converted into electrical energy by appropriate devices.
• The various forms of energy are interconnected; under suitable conditions they transform from one form into another. Thermodynamics studies and quantifies these transformations.

Note: Thermodynamics is the science that studies the combined effects of heat and work on changes in the state of matter, constrained by the laws of thermodynamics.

Why Study Thermodynamics?

• Feasibility of a process: Provides criteria to predict whether a process (including chemical reactions) is feasible or spontaneous under given conditions.
• Extent of a process: Helps determine to what extent a process can proceed before equilibrium is attained.
• Efficiency of a process: Allows assessment of how effectively energy can be converted from one form to another using the laws (zeroth, first, second and third) of thermodynamics.

Note: The laws of thermodynamics apply when a system is in equilibrium or moves from one equilibrium state to another.

System, Surroundings and BoundarySystem

A thermodynamic system is the portion of the universe chosen for study. It may be a fixed amount of matter or a region in space. A system can exchange mass, energy, or both with its surroundings.

System, Surrounding, and Boundary

Surroundings

The surroundings are everything external to the system. The system together with its surroundings constitutes the universe (in thermodynamic terms). The system may exchange heat with the surroundings and do work on the surroundings.

Boundary

The boundary separates the system and the surroundings. It may be real or imaginary, rigid or flexible, and either a conductor or an insulator of heat.

System + Boundary + Surroundings = Universe

• It may be diathermic (able to conduct heat or pass heat freely)or adiabatic (there is no heat transfer from in or out of the system).
  
  Difference in Diathermic and Adiabatic Wall
• The boundary may be real or imaginary; it may be rigid or non-rigid; it may be a conductor or a non-conductor of heat.

System Thermodynamics

Types of Thermodynamic SystemOpen system

• An open system can exchange both matter and energy with the surroundings. Example: a beaker of liquid open to the atmosphere.

Open System

Closed system

• A closed system can exchange only energy (heat or work) with the surroundings but not matter. Example: a sealed vessel containing liquid.

Closed System

Isolated system

• An isolated system exchanges neither matter nor energy with the surroundings. Example: an ideally insulated flask (a thermos flask approximates this).

Isolated System

Basic Thermodynamic Terms: Internal Energy, Work and Heat

Chemical thermodynamics relates work, heat and chemical reactions or physical changes. The principal quantities used are internal energy (U), work (w) and heat (q). The sections below define and explain each.

(a) Internal Energy

Internal energy refers to the total energy contained within the system. It includes all microscopic forms of energy such as translational, rotational and vibrational kinetic energies, potential energies from intermolecular forces, chemical energy, electronic energy and others.

• Internal energy is the total microscopic energy of the system.
• It is the sum of kinetic and potential energies of all particles: translational, rotational, vibrational, electronic, chemical, nuclear, etc.

Represented symbolically: U = translational KE + rotational KE + vibrational KE + chemical energy + electronic energy + nuclear energy + potential energy

• The internal energy U of a system may change when heat passes into or out of the system, when work is done on or by the system, or when matter enters or leaves the system.
• Internal energy is an extensive property and a state function.
• The absolute value of U cannot be measured directly, but changes in internal energy ΔU are experimentally determinable.

Change in internal energy:

ΔU = U₂ - U₁

First law (for a closed system, sign convention where w is work done on the system):

ΔU = q + w

In many cases, exothermic processes decrease internal energy (ΔU < 0), while endothermic processes increase it (ΔU > 0).

U depends on temperature, pressure, volume and the amount of substance.

Factors affecting internal energy: Internal energy changes when heat is transferred, when work is done on/by the system, or when matter enters or leaves the system.

(b) Work

Work is the energy transferred between a system and its surroundings by any process other than heat transfer. Mechanical work is common in thermodynamics and is governed by external forces like pressure.

• If an object is displaced through a distance dx against a force F, the elementary work done is w = F × dx.
• Work associated with volume change against an external pressure is called pressure-volume work or mechanical work.

Mechanical (pressure-volume) work:

w = -P_ext × (V₂ - V₁) = -P_ext ΔV

• Work (w) is a path-dependent function (a path function).
• Work done on a system increases the system's energy; work done by the system decreases the system's energy.

Sign convention for work:

Work done on the system ⇒ w is positive.

Work done by the system ⇒ w is negative.

Expansion ⇒ system does work on surroundings ⇒ w is negative.

Compression ⇒ work is done on system ⇒ w is positive.

(c) Heat

Heat (q) is the energy transferred between a system and its surroundings because of a temperature difference. Heat flows spontaneously from higher temperature to lower temperature.

Change in internal energy formula

• Heat transfer occurs by conduction, convection or radiation depending on the system and boundary conditions.
• Heat and thermodynamics together allow engineers and scientists to design processes that harness chemical and thermal energy efficiently.

The Laws of Thermodynamics

Important: The laws apply when a system is in equilibrium or moves between equilibrium states.

Laws of Thermodynamics

• Zeroth law of thermodynamics: If two bodies are each in thermal equilibrium with a third body, then they are in thermal equilibrium with each other.
• First law of thermodynamics: Energy is conserved. For a closed system, the change in internal energy equals heat added to the system plus work done on the system: ΔU = q + w.
• Second law of thermodynamics: The entropy of the universe (considered as an isolated system) tends to increase; natural processes have a preferred direction and are not all reversible.
• Third law of thermodynamics: The entropy of a perfect crystal at absolute zero temperature is zero.

State and Path Functions

A property whose value does not depend on the path taken to reach that value is called a state function (or point function). Its value depends only on the current state described by macroscopic properties.

• When macroscopic properties of a system have definite values, the system is in a definite state.
• Changing any macroscopic property changes the state of the system.

Thus the state of a system is fixed by its macroscopic properties (state variables).

State Variables (State Functions)

State or condition of a system is described by certain measurable properties & these measurable properties are called state variables.

State variables or the thermodynamics parameters depend only upon the initial and final states of the system and are independent of the manner in how the change is brought are called state functions.

Note: A state function depends only on the initial and final states of the system; it does not depend on the path followed.

Common state functions include: internal energy (U), enthalpy (H), entropy (S), Gibbs free energy (G), pressure (P), temperature (T), volume (V), mass, etc.

1. Pressure: Pressure is a measure of the average force exerted by the constituent molecules per unit area on the container walls. Pressure does not depend on the path of the molecules and thus it is a state function.
2. Temperature: Temperature is defined as the measure of the average kinetic energy of the atoms or molecules in the system. Temperature measures a property of a state of a system, irrespective of how it got there and thus it is a state function.
3. Volume: Volume is the amount of physical space occupied by a substance and it will not be dependent on the path followed. Thus, the volume is a state function
4. Mass: The measure of the amount of matter in an object is known as mass and is usually measured in grams (g) or kilograms (kg). Mass measures the quantity of matter regardless of both its location in the universe and the gravitational force applied to it and thus it is a state function.
5. Internal energy: It can be defined as the sum of all kind of energy associated with molecular motions.
   The internal energy of ideal gases is a function of temperature only (Joule’s law) and for ideal gases, internal energy depends only on temperature (Joule’s law). For real gases, it depends on temperature and volume. So it can be seen that since internal energy depends on quantities like P, T, V which are state functions, the internal energy is also a state function.
6. Gibbs free energy: The enthalpy of the system at any point minus the product of the temperature times the entropy of the system is Gibbs free energy of the system
   G = H – TS
   The Gibbs free energy of the system is a state function because it is defined in terms of thermodynamic properties that are state functions.
7. Entropy: Entropy is a measure of the disorder or randomness of a system and it’s totally independent of the path through which the system has achieved that state also it’s unique to the current state of the system.

Path Functions

Path functions depend on the path taken between initial and final states. Heat (q) and work (w) are the primary examples of path functions in classical thermodynamics.

Example: Heat and work are path functions because their values depend on how the process is carried out, not only on the end states.

Intensive and Extensive Properties

Intensive and extensive properties are terms used to describe physical properties. The terms were described by Richard C. Tolman in 1917.

Intensive Properties

• Intensive properties do not depend on the amount of substance present. Examples: temperature, density, colour, boiling point, refractive index, hardness, melting point, odour, luster.
• These properties remain the same for a substance regardless of sample size.

Extensive Properties

• Extensive properties depend on the amount of substance present. Examples: mass, volume, size, weight, length, energy.
• Extensive properties are additive for subsystems; e.g., total mass or total energy of two subsystems is the sum of their masses/energies.

The ratio of two extensive properties is often an intensive property. For example, density = mass/volume (both extensive) is intensive.

Way to tell them apart: Combine two identical samples. If the property doubles, it is extensive. If it remains the same, it is intensive.

Understanding whether a property is intensive or extensive is crucial when characterising and identifying substances.

• The two types of physical properties of matter are intensive and extensive.
• Intensive properties do not depend on the quantity; examples include density, state of matter, temperature.
• Extensive properties depend on sample size; examples include volume, mass, and energy.

Reversible and Irreversible ProcessesReversible process

• A reversible process proceeds through an infinite number of infinitesimal steps such that the system is in equilibrium at each step. The process can be reversed by an infinitesimal change in a variable, restoring both system and surroundings to their original states.
• In a reversible process there is no net loss of useful energy and the maximum possible work is obtained for the change between two states.
• Reversible processes are idealised and proceed infinitely slowly; they are theoretical limits for efficiency of engines and processes.
• For reversible expansion/compression, Pext = Pint (external pressure equals internal pressure at each step).

Reversible & Irreversible Processes

Irreversible process

An irreversible process is one that cannot be exactly reversed to restore both the system and surroundings to their initial states. Most natural processes are irreversible.

• Irreversible processes are real processes and typically occur at finite rates (not infinitely slowly).
• For irreversible processes Pext ≠ Pint.
• Irreversible processes involve dissipative effects such as friction, turbulence, spontaneous heat flow across finite temperature differences, and unrestrained expansion.

Comparison between reversible and irreversible processes

Some Important Questions-

Q.1. Give examples of state functions.

Ans: Examples of state functions are:

• Pressure
• Mass
• Temperature
• Enthalpy
• Internal energy
• Entropy
• Volume

Q.2. A Thermos flask is a ______

1. Diathermic system
2. Only a closed system
3. Isolated system
4. Only an adiabatic system

Ans: Isolated System
Solution: A thermos flask is an isolated system. It is an insulating storage vessel that keeps a hotter or cooler item at its respective temperature for a long time, approximating no exchange of heat or matter with the surroundings.

Thermos Flask

Q.3. In thermodynamics, heat, and work are :

1. Point functions
2. Extensive thermodynamic state variables
3. Path functions
4. Intensive thermodynamic state variables

Ans: Heat and work are Path Functions

Solution:

• Path functions depend on the path taken by the thermodynamic process from initial to final state. Their differentials are inexact.
• Heat and work are the primary examples of path functions.

Q.4. Spot the odd one out

1. Specific enthalpy
2. Kinetic energy
3. Work
4. Pressure

Answer: Option 3: work
Solution: Work is a path function (depends on the path), while specific enthalpy, kinetic energy (as a property for a system), and pressure are state functions.

Q.5. Which of the following is an intensive property of a thermodynamic system?

1. Temperature
2. Mass
3. Energy
4. Volume

Answer: Option 1: temperature
Solution: Temperature is intensive (does not depend on the amount of substance). Mass, energy and volume are extensive properties.