Thermodynamic Property: Definition, Relations, Calculator, Examples

Thermodynamic Property

Edited By Shivani Poonia | Updated on Jul 02, 2025 06:31 PM IST

We learn that thermodynamics is the branch of Physics and Chemistry that deals with heat, temperature, work, and their relation to energy, radiation, and the physical properties of matter. Thermodynamics deals with large systems. We can see the concept of thermodynamics applied in our everyday lives in cars, bikes, ACs, thermal flasks, refrigerators, washing machines, and many more.

This Story also Contains

Properties of System
2. Extensive Properties
Some Solved Examples
Summary

Thermodynamic Property

There are plenty of characteristics and concepts in thermodynamics that help in defining the system. Multiple external and internal properties contribute to the result in terms of performance and reflect back on the total energy of the system. There exist many measurable properties of the thermodynamic system that help describe the current physical state of the system. These properties define the characteristic features of a system. Thermodynamic properties can be applied to energy systems and thermal or nuclear power plants. This article allows us to understand in detail the thermodynamic properties of the system.

Properties of System

All macroscopic properties of a system irrespective of the fact whether they are state variables or not are divided into two types:

1. Intensive Properties

Such properties remain the same on any division in the system that is, do not depend upon the amount of substance present in the system.

Example: Temperature, pressure, concentration, density, viscosity, surface tension, specific heat, refractive index, pH, EMF of the dry cell, vapor pressure, dipole-moment etc.

2. Extensive Properties

Such properties depend upon the amount of substance that is, their values are different in the divided system than in the entire system

Examples, Mass, volume, energy, work, internal energy, entropy, enthalpy, heat capacity, and length.

An extensive property can be made intensive specifying it in unit amount of matter.

Example, Density ≈( Mass Volume )

Some Solved Examples

Example 1: Which of the following is an extensive property?

1) Enthalpy

2)Concentration

3)Density

4)Viscosity

Solution

Extensive Properties - Such properties depend upon the amount of substance that is, their values are different in the divided system than in the entire system Examples are mass, volume, energy, work, internal energy, entropy, enthalpy, heat capacity, and length. An extensive property can be made intensive specifying its In-unit amount of matter.

Example, Density ≈( Mass Volume )

Therefore we can say enthalpy depends upon the quantity of substance.

Hence, the correct answer is the option (1).

Example 2: Identify intensive property from the following

1)Volume

2)Mass

3)Enthalpy

4) Temperature

Solution

The properties that do not depend upon the quantity of matter present in the system or the size of the system are called intensive properties. For example temperature, pressure, concentration, density, dipole moment, refractive index, etc. Therefore we can say the temperature is independent of the mass of the substance.

Hence, the correct answer is option (4).

Example 3: Entropy of a pure substance is an ____________

1)point function

2)extensive property

3)intensive property

4)none of the mentioned

Solution

The entropy of a substance is an extensive property as it depends upon the amount of system taken.
Hence, the correct answer is the option (2).

Example 4: Which among the following is extensive property?

1)density

2)temperature

3) enthalpy

4)specific heat capacity

Solution

Extensive Property - Properties of the system depend on mass or the size of the system. All types of energy, mass, length, total heat capacity. Enthalpy depends on the amount of the system.

Hence, the answer is the option (3).

Example 5: The intensive property among these quantities is

1)enthalpy

2) mass/volume

3)mass

4)Volume

Solution

Intensive Property - Properties of the system which does not depend on mass or the size of the system. Pressure, Temperature, Density, Molar heat capacity, Specific heat capacity, concentration, Dipole moment, Viscosity, Refractive index, pH, Surface tension, EMF of Dry cell, Specific gravity, R. Mass/Volume is independent of the amount of the system.

Hence, the answer is the option (2).

Summary

Temperature measures the average kinetic energy of particles in a substance and is essential for understanding heat transfer and energy changes. It’s commonly measured in Kelvin (K), Celsius (°C), or Fahrenheit (°F), with the Kelvin scale being crucial in scientific contexts. Pressure, the force per unit area exerted by gas molecules, is vital for describing the states of gases and liquids and is measured in units like Pascal (Pa) or atmosphere (atm). It plays a key role in the ideal gas law linking pressure, volume, temperature, and the number of molecules. Temperature measures how hot or cold something is and reflects the average kinetic energy of particles in a substance. It is fundamental in thermodynamics because it influences heat transfer and energy changes. Common units of temperature include Kelvin (K), Celsius (°C), and Fahrenheit (°F). The Kelvin scale is significant in scientific contexts, with absolute zero (0 K) being the point where particles have minimum kinetic energy.

Frequently Asked Questions (FAQs)

1. How do intensive and extensive properties differ in thermodynamics?

Intensive properties are independent of the system's size or amount of material, such as temperature or pressure. Extensive properties depend on the system's size, like volume or total energy. Dividing an extensive property by the amount of substance gives a corresponding intensive property.

2. Why is entropy considered a state function?

Entropy is a state function because its value depends only on the current state of the system, not on its history or how it reached that state. This means the change in entropy between two states is independent of the path taken between them.

3. What is the significance of Gibbs free energy in chemical reactions?

Gibbs free energy is crucial for determining the spontaneity of chemical reactions at constant temperature and pressure. A negative change in Gibbs free energy indicates a spontaneous process, while a positive change indicates a non-spontaneous one. It combines enthalpy and entropy effects into a single value.

4. How does the First Law of Thermodynamics relate to energy conservation?

The First Law of Thermodynamics is essentially the law of energy conservation applied to thermodynamic systems. It states that the total energy of an isolated system is constant; energy can be transformed from one form to another, but it cannot be created or destroyed.

5. How does the concept of free energy relate to chemical equilibrium?

At chemical equilibrium, the Gibbs free energy of the system is at a minimum. The change in Gibbs free energy (ΔG) approaches zero as the reaction nears equilibrium. This concept helps predict the direction of spontaneous change and the point at which a reaction will stop.

6. What is a thermodynamic property?

A thermodynamic property is a measurable characteristic of a system that depends on its current state, not on how it reached that state. These properties help describe the system's behavior and can be used to predict changes in energy, work, and heat transfer.

7. How does pressure affect the boiling point of a liquid?

Increasing pressure raises the boiling point of a liquid. This is because higher pressure makes it harder for vapor bubbles to form and escape the liquid surface. Conversely, decreasing pressure lowers the boiling point, which is why water boils at a lower temperature at high altitudes.

8. What is the difference between heat and temperature?

Heat is a form of energy transfer between systems due to temperature differences, while temperature is a measure of the average kinetic energy of particles in a substance. Heat flows from higher to lower temperatures, but temperature is a property of a system regardless of energy transfer.

9. What is the physical meaning of enthalpy?

Enthalpy is a measure of the total heat content of a system. It represents the sum of the system's internal energy and the product of its pressure and volume. Changes in enthalpy indicate the amount of heat absorbed or released during a process at constant pressure.

10. What is the significance of the triple point in a phase diagram?

The triple point is the unique combination of temperature and pressure at which a substance can exist simultaneously in solid, liquid, and gas phases in equilibrium. It's a fixed point used in defining temperature scales and understanding phase behavior.

11. Why are some processes irreversible in thermodynamics?

Processes are irreversible when they increase the entropy of the universe. In reality, all natural processes involve some degree of irreversibility due to factors like friction, heat transfer across finite temperature differences, and the spontaneous mixing of substances. Perfectly reversible processes are idealized concepts.

12. How does the Second Law of Thermodynamics explain the direction of spontaneous processes?

The Second Law states that the total entropy of an isolated system always increases for spontaneous processes. This law explains why heat flows from hot to cold objects, why chemical reactions proceed in certain directions, and why perfect efficiency in heat engines is impossible.

13. Why is the efficiency of a Carnot engine considered the maximum possible?

The Carnot engine is a theoretical heat engine operating in a perfectly reversible cycle between two temperature reservoirs. Its efficiency represents the maximum possible for any heat engine working between those temperatures. Real engines are always less efficient due to irreversible processes and heat losses.

14. How does the van der Waals equation improve upon the ideal gas law?

The van der Waals equation accounts for two factors ignored by the ideal gas law: the finite size of gas molecules and the attractive forces between them. It introduces two constants: one related to molecular volume and another to intermolecular attractions. This equation provides a more accurate description of real gas behavior, especially at high pressures and low temperatures.

15. How does the concept of partial pressure relate to Raoult's law and Henry's law?

Partial pressure is the pressure exerted by a single component in a gas mixture. Raoult's law relates the vapor pressure of an ideal solution to the vapor pressure of pure components and their mole fractions. Henry's law describes the solubility of gases in liquids at low concentrations. Both laws use partial pressures to describe gas-liquid equilibria.

16. How does the concept of chemical affinity relate to the driving force of chemical reactions?

Chemical affinity is the tendency of chemical species to combine with or separate from one another. It's directly related to the change in Gibbs free energy of a reaction. A higher affinity indicates a greater driving force for the reaction to proceed. This concept is fundamental in understanding reaction spontaneity and equilibrium.

17. How does the concept of activity relate to chemical equilibrium constants?

Activity is an effective concentration that accounts for non-ideal behavior in solutions or gas mixtures. In chemical equilibrium, the equilibrium constant is properly expressed in terms of activities rather than concentrations or partial pressures. This concept allows for more accurate predictions of equilibrium compositions in non-ideal systems.

18. What is the relationship between internal energy and temperature for an ideal gas?

For an ideal gas, the internal energy depends only on temperature, not on pressure or volume. This is because ideal gas particles are assumed to have no intermolecular forces, so their energy is purely kinetic and directly related to temperature.

19. How does the concept of chemical potential relate to phase equilibrium?

Chemical potential is the driving force for mass transfer between phases. At phase equilibrium, the chemical potential of a substance is the same in all coexisting phases. This concept helps explain phenomena like vapor pressure and the distribution of solutes between phases.

20. What is the physical interpretation of absolute zero temperature?

Absolute zero (0 K or -273.15°C) is the theoretical temperature at which all molecular motion ceases. It's the lowest possible temperature, where a system would have minimum thermal energy. However, due to quantum effects, true absolute zero can never be reached.

21. How does enthalpy of formation relate to bond energies?

The enthalpy of formation is related to bond energies but also includes intermolecular forces. It represents the energy change when forming a compound from its elements in their standard states. This value can be calculated by considering the energy required to break bonds in reactants and the energy released when forming new bonds in products.

22. How does Le Chatelier's principle relate to thermodynamic equilibrium?

Le Chatelier's principle states that when a system at equilibrium is disturbed, it will shift to counteract the change. This principle is a consequence of the system trying to minimize its Gibbs free energy. Temperature, pressure, or concentration changes cause the system to re-establish equilibrium by shifting the reaction.

23. What is the significance of the Joule-Thomson effect in thermodynamics?

The Joule-Thomson effect describes the temperature change of a gas when it expands through a valve or porous plug without doing external work. This effect is crucial in processes like gas liquefaction and is used in refrigeration cycles. The sign of the temperature change depends on the gas and its initial conditions.

24. How does the concept of fugacity extend the ideal gas law to real gases?

Fugacity is a measure of the tendency of a substance to escape from a phase. It's used to describe the behavior of real gases, extending the concept of pressure from ideal gas law. For an ideal gas, fugacity equals pressure, but for real gases, it accounts for intermolecular forces and deviations from ideal behavior.

25. What is the physical meaning of partial molar quantities in thermodynamics?

Partial molar quantities represent how an extensive property of a solution changes when the amount of one component is varied while keeping temperature, pressure, and the amounts of other components constant. They are crucial for understanding the behavior of mixtures and solutions.

26. How does the Maxwell-Boltzmann distribution relate to temperature and molecular speeds?

The Maxwell-Boltzmann distribution describes the range of molecular speeds in a gas at a given temperature. As temperature increases, the distribution broadens and shifts to higher speeds, reflecting increased molecular kinetic energy. This concept links microscopic particle behavior to macroscopic temperature.

27. What is the significance of the critical point in a phase diagram?

The critical point is the end point of the liquid-vapor coexistence curve in a phase diagram. Beyond this point, the distinction between liquid and gas phases disappears, forming a supercritical fluid. Understanding the critical point is crucial for processes involving supercritical fluids and for designing efficient separation processes.

28. How does the concept of activity coefficient account for non-ideal behavior in solutions?

The activity coefficient is a factor used to account for deviations from ideal behavior in solutions. It modifies the concentration term in thermodynamic equations to reflect the effective concentration of a species in a non-ideal mixture. This concept is crucial for accurate predictions of chemical equilibria and phase behavior in real solutions.

29. What is the thermodynamic basis for the Nernst equation in electrochemistry?

The Nernst equation relates the reduction potential of an electrochemical reaction to the standard electrode potential and the activities of the species involved. It's derived from the relationship between Gibbs free energy and cell potential, providing a link between thermodynamics and electrochemistry.

30. What is the significance of the partition function in statistical thermodynamics?

The partition function is a central concept in statistical thermodynamics that connects microscopic properties of molecules to macroscopic thermodynamic properties. It sums over all possible energy states of a system, allowing the calculation of various thermodynamic quantities like internal energy, entropy, and heat capacity.

31. How does the concept of residual properties help in understanding real gas behavior?

Residual properties are the differences between the actual properties of a real gas and those of an ideal gas at the same temperature and pressure. They provide a measure of the deviation from ideal behavior and are useful in calculating thermodynamic properties of real gases using equations of state.

32. What is the thermodynamic explanation for the Mpemba effect?

The Mpemba effect is the counterintuitive phenomenon where, under certain conditions, hot water can freeze faster than cold water. While still debated, thermodynamic explanations involve factors like evaporation, convection currents, dissolved gases, and supercooling. It highlights the complexity of non-equilibrium thermodynamic processes.

33. How does the concept of exergy differ from energy in thermodynamic analysis?

Exergy is the maximum useful work that can be extracted from a system as it reaches equilibrium with its surroundings. Unlike energy, which is conserved, exergy can be destroyed in irreversible processes. This concept is crucial in assessing the efficiency and sustainability of energy conversion processes.

34. What is the thermodynamic basis for the Clausius-Clapeyron equation?

The Clausius-Clapeyron equation relates the pressure and temperature at which two phases of a substance coexist in equilibrium. It's derived from the condition that the chemical potentials of the two phases must be equal at equilibrium. This equation is crucial for understanding vapor pressure curves and phase transitions.

35. How does the Third Law of Thermodynamics relate to absolute entropy?

The Third Law states that the entropy of a perfect crystal at absolute zero is zero. This provides a reference point for calculating absolute entropies of substances at other temperatures. It's crucial for understanding low-temperature behavior of materials and for calculating standard entropies of formation.

36. What is the significance of the Helmholtz free energy in thermodynamics?

Helmholtz free energy (A = U - TS) is a thermodynamic potential that's minimized at equilibrium for systems at constant volume and temperature. It's particularly useful in studying phase transitions and chemical reactions in closed systems where volume, not pressure, is held constant.

37. How does the concept of fugacity coefficient relate to the compressibility factor?

The fugacity coefficient is the ratio of fugacity to pressure and is related to the compressibility factor (Z) through an integral of (Z-1)/P with respect to pressure. It provides a measure of the deviation of a real gas from ideal gas behavior and is crucial in high-pressure thermodynamics and phase equilibria calculations.

38. What is the thermodynamic explanation for the Leidenfrost effect?

The Leidenfrost effect occurs when a liquid comes into contact with a surface significantly hotter than its boiling point. A vapor layer insulates the bulk liquid from the hot surface. Thermodynamically, this involves rapid vaporization, creating a high-pressure region that supports the droplet, minimizing heat transfer and prolonging the liquid's lifetime.

39. How does the concept of excess properties help in understanding non-ideal mixtures?

Excess properties are the differences between the actual properties of a mixture and those of an ideal mixture at the same temperature, pressure, and composition. They quantify the non-ideal behavior of mixtures due to molecular interactions. Understanding excess properties is crucial for accurate predictions of mixture behavior in chemical engineering and materials science.

40. What is the thermodynamic basis for the Gibbs-Duhem equation?

The Gibbs-Duhem equation relates changes in chemical potentials of components in a system. It's derived from the fundamental equation of thermodynamics and the definition of Gibbs free energy. This equation is crucial in phase equilibria calculations and in understanding the interdependence of intensive variables in multicomponent systems.

41. What is the significance of the Joule-Thomson coefficient in gas expansion processes?

The Joule-Thomson coefficient describes how the temperature of a gas changes with pressure during a constant enthalpy expansion. It's crucial in understanding gas behavior during throttling processes, such as in refrigeration cycles. The sign of this coefficient determines whether a gas cools or heats upon expansion.

42. What is the thermodynamic explanation for the heat of mixing in solutions?

The heat of mixing arises from the energy changes when different substances are combined to form a solution. It can be exothermic or endothermic, depending on whether the interactions between unlike molecules are more or less favorable than those between like molecules. This concept is crucial in understanding solution thermodynamics and non-ideal mixing behavior.

43. How does the concept of fugacity apply to phase equilibria in multicomponent systems?

In multicomponent systems, the fugacity of each component must be equal in all phases at equilibrium. This principle extends the concept of vapor-liquid equilibrium to more complex systems. It's crucial in designing separation processes, understanding phase behavior of mixtures, and predicting equilibrium compositions in various industrial processes.

44. What is the significance of the Maxwell relations in thermodynamics?

Maxwell relations are equations that relate partial derivatives of thermodynamic quantities. They are derived from the mathematical properties of exact differentials and the definitions of thermodynamic potentials. These relations are powerful tools for deriving useful equations and calculating thermodynamic properties that are difficult to measure directly.

45. How does the concept of osmotic pressure relate to chemical potential?

Osmotic pressure is the pressure required to prevent the flow of solvent across a semipermeable membrane into a solution. It's directly related to the difference in chemical potential of the solvent between the pure solvent and the solution. This concept is crucial in understanding membrane processes, biological systems, and colligative properties.

46. What is the thermodynamic basis for the Gibbs phase rule?

The Gibbs phase rule relates the number of degrees of freedom in a thermodynamic system to the number of components and phases present. It's derived from the condition of phase equilibrium and the Gibbs-Duhem equation. This rule is fundamental in understanding phase diagrams and predicting the behavior of multiphase, multicomponent systems.

47. What is the significance of the Trouton's rule in vaporization processes?

Trouton's rule states that the molar entropy of vaporization is approximately constant (about 88 J/mol·K) for many liquids at their normal boiling points. This empirical rule provides a quick estimate of enthalpies of vaporization and is useful in understanding the behavior of liquids and their intermolecular forces.

48. How does the concept of enthalpy-entropy compensation relate to chemical and biochemical processes?

Enthalpy-entropy compensation refers to the phenomenon where changes in enthalpy and entropy often offset each other in chemical and biochemical processes. This concept is important in understanding the driving forces behind reactions, especially in biological systems, where small changes in Gibbs free energy can result from large, compensating changes in enthalpy and entropy.

49. What is the thermodynamic explanation for the Donnan equilibrium in biological systems?

The Donnan equilibrium describes the unequal distribution of ions across a semipermeable membrane when one ion cannot pass through the membrane.