Thermodynamics for NEET 2027: Notes, Formulas and Important Concepts
Mishika Gupta17 Oct, 2026
Thermodynamics can be challenging because it involves several concepts, sign conventions and numerical calculations. Understanding the relationship between heat, work and internal energy is essential to solving problems correctly. Confusion about formulas, enthalpy changes or the conditions for spontaneity can make it difficult to apply these concepts in NEET Chemistry questions.
These notes for NEET 2027 cover the chapter’s important concepts, formulas and examples to help you revise effectively. You can review thermodynamic systems, pressure-volume work, the first law, enthalpy, entropy and Gibbs free energy while learning how to apply the relevant equations to numerical problems.
Thermodynamics Introduction
Thermodynamics is the study of energy changes and energy transfer during physical and chemical processes. It explains how a system exchanges energy with its surroundings through heat and work.
For example, a marker held at a height has gravitational potential energy. When it falls, its potential energy decreases and is converted into other forms of energy, such as kinetic energy. The energy is transferred or transformed rather than destroyed.
Meaning and Scope of Thermodynamics
The word thermodynamics comes from two terms:
-
Thermo: Heat or energy.
-
Dynamics: Change or movement.
Thermodynamics helps you understand:
-
How energy changes during physical and chemical processes.
-
How heat and work transfer energy between a system and its surroundings.
-
How internal energy and enthalpy change during reactions.
-
How entropy and Gibbs free energy help determine spontaneity.
System, Surroundings and Boundary
A system is the part of the universe selected for observation. It may be a reaction mixture, a gas inside a cylinder or a liquid in a container.
The surroundings include everything outside the system, while the boundary separates the system from its surroundings.
[\text{Universe}=\text{System}+\text{Surroundings}]
Energy can cross the boundary as heat or work, depending on the type of system and process.
Types of Thermodynamic Systems
|
System type |
Matter exchange |
Energy exchange |
Example |
|
Open system |
Yes |
Yes |
An open cup of hot tea |
|
Closed system |
No |
Yes |
A sealed vessel containing hot tea |
|
Isolated system |
No |
No |
An ideal thermos flask |
An open cup allows heat transfer and the escape of water vapour. A sealed vessel can exchange energy with its surroundings but not matter. An isolated system exchanges neither matter nor energy under the assumed conditions.
Heat and Work: Modes of Energy Transfer
Heat and work describe how energy moves between a system and its surroundings. They are not state properties of the system.
Examples of energy forms include:
-
Chemical energy
-
Mechanical energy
-
Rotational energy
-
Vibrational energy
Heat transfer occurs because of a temperature difference. Work occurs when energy is transferred through an organised process, such as the expansion or compression of a gas against an external pressure.
Pressure-Volume Work
Pressure-volume work occurs when a gas expands or contracts against an external pressure. Consider a gas inside a cylinder fitted with a movable piston. The gas particles collide with the piston and exert a force. Pressure is defined as:
[P=\frac{F}{A}]
where (F) is force and (A) is area.
The piston moves when the forces acting on it are unbalanced. Expansion occurs when the gas pushes the piston outward, while compression occurs when the surroundings push it inward.
Pressure-Volume Work Formula
At constant external pressure: [\boxed{w=-P_{\mathrm{ext}}\Delta V}]
where:[\Delta V=V_f-V_i]
Here, (V_f) is the final volume and (V_i) is the initial volume.
|
Condition |
Volume change |
Sign of work |
|
Expansion |
(\Delta V>0) |
(w<0) |
|
Compression |
(\Delta V<0) |
(w>0) |
|
No volume change |
(\Delta V=0) |
(w=0) |
Remember: In the chemistry sign convention, expansion work is negative because the system performs work on its surroundings. Compression work is positive because work is done on the system.
Work Units and Conversion
The SI unit of work is the joule (J).
Useful conversions include:
-
(1\ \mathrm{L,atm}\approx101.3\ \mathrm{J})
-
(1\ \mathrm{bar}=10^5\ \mathrm{Pa})
-
(1\ \mathrm{Pa,m^3}=1\ \mathrm{J})
Keep pressure and volume units consistent when solving numerical problems.
Reversible and Irreversible Processes
A reversible process is an idealised process that proceeds through a series of infinitesimally small changes and can be reversed by an infinitesimal change in conditions. An irreversible process cannot be reversed in this way without leaving changes in the system or surroundings.
|
Feature |
Reversible process |
Irreversible process |
|
Nature |
Idealised |
Common in real processes |
|
Progress |
Through infinitesimal changes |
Through finite changes |
|
Driving force |
Infinitesimal difference from equilibrium |
Finite difference from equilibrium |
|
Reversal |
Can be reversed through infinitesimal changes |
Cannot be reversed through infinitesimal changes |
A reversible process is not simply any slow process; it must also avoid dissipative effects such as friction.
Pressure-Volume Curves and Cyclic Processes
A pressure-volume (P–V) diagram shows how pressure and volume change during a process. For a reversible process, the work done by the system is:
[w=-\int_{V_i}^{V_f}P,dV]
The area under the curve represents the magnitude of pressure-volume work. Its sign depends on whether the process involves expansion or compression. A cyclic process returns the system to its initial state, forming a closed loop on a P–V diagram.
For a complete cycle:[\Delta U=0]
The net work corresponds to the signed area enclosed by the loop. Under the chemistry sign convention, a clockwise cycle generally has negative net work, while an anticlockwise cycle has positive net work.
Important Thermodynamic Processes
Thermodynamic processes are classified according to the property that remains constant or the condition that applies.
|
Process |
Condition |
Key relationship |
|
Isothermal |
Temperature remains constant |
(\Delta T=0) |
|
Isobaric |
Pressure remains constant |
(P=\text{constant}) |
|
Adiabatic |
No heat exchange |
(q=0) |
|
Isochoric |
Volume remains constant |
(\Delta V=0,\ w=0) |
Types of Thermodynamic Processes
Thermodynamic processes are classified based on changes in temperature, pressure, volume, and heat exchange, with each type following specific conditions and equations.
|
Process |
Definition |
Key Condition |
Work Done (w) |
Important Relation |
|
Isothermal Process |
A process that occurs at constant temperature. For an ideal gas, internal energy depends only on temperature. |
T=0, so U=0 |
For reversible expansion or compression: \(\displaystyle w=-nRT\ln\frac{V_2}{V_1}\) |
\(\displaystyle w=-2.303nRT\log\frac{V_2}{V_1}\) \(\displaystyle w=-nRT\ln\frac{P_1}{P_2}\) |
|
Isobaric Process |
A process that occurs at constant pressure. |
P=0 |
\(\displaystyle w=-P\Delta V\) |
For an ideal gas at constant pressure: \(\displaystyle w=-nR\Delta T\) |
|
Adiabatic Process |
A process in which no heat is exchanged between the system and surroundings. |
q=0 |
From the first law, \(\displaystyle \Delta U=w\) |
\(\displaystyle \Delta U=q+w=w\) |
|
Isochoric Process |
A process that occurs at constant volume. |
V=0 |
\(\displaystyle w=0\) |
\(\displaystyle \Delta U=q_v\) |
Important points to remember:
-
For isothermal processes, use temperature in kelvin and choose R according to the units.
-
For isobaric processes, w=-nRT applies to an ideal gas when pressure remains constant, and the external pressure equals the gas pressure.
-
For an adiabatic process, U=w follows from the first law using the chemistry sign convention.
-
For an isochoric process, qv represents the heat exchanged at constant volume.
Internal Energy and the First Law of Thermodynamics
Internal energy ((U)) is the total microscopic energy of a system, including contributions from molecular motion and interactions.
|
Concept |
Formula or Rule |
Explanation |
|
Change in Internal Energy |
U=Uf-Ui |
Difference between final and initial internal energy. |
|
First Law of Thermodynamics |
U=q+w |
Energy is conserved; internal energy changes due to heat and work. |
|
Positive Heat (q>0) |
q>0 |
Heat is absorbed by the system. |
|
Negative Heat (q<0) |
q<0 |
Heat is released by the system. |
|
Positive Work (w>0) |
w>0 |
Work is done on the system. |
|
Negative Work (w<0) |
w<0 |
Work is done by the system. |
|
State Functions |
U,H,P,V,T,S |
Depend only on the initial and final states, not on the path followed. |
|
Path Functions |
q,w |
Depend on the path taken during the process. |
|
Isothermal Process for an Ideal Gas |
\(\Delta T=0 \implies \Delta U=0\) |
Internal energy remains constant because it depends only on temperature for an ideal gas. |
Note: The formulas use the chemistry sign convention, where U=q+w.
Thermochemistry and Enthalpy
Thermochemistry studies heat changes associated with chemical reactions and physical transformations.
|
Concept |
Formula |
Explanation |
|
Enthalpy |
\(\displaystyle H=U+PV\) |
Sum of internal energy and the product of pressure and volume. |
|
Change in Enthalpy |
\(\displaystyle \Delta H=\Delta U+\Delta(PV)\) |
Relates the changes in enthalpy, internal energy, pressure and volume. |
|
Enthalpy–Internal Energy Relation |
\(\displaystyle \Delta H=\Delta U+\Delta n_gRT\) |
Applies to reactions involving ideal gases under the usual conditions. |
|
Change in Gaseous Moles |
\(\displaystyle \Delta n_g=\text{Moles of gaseous products}-\text{Moles of gaseous reactants}\) |
Count only gaseous species, using their stoichiometric coefficients. |
|
Heat at Constant Volume |
\(\displaystyle q_v=\Delta U\) |
Applies when pressure-volume work is the only work involved. |
|
Heat at Constant Pressure |
\(\displaystyle q_p=\Delta H\) |
Applies when pressure-volume work is the only work involved. |
Important Enthalpy Terms
|
Enthalpy Term |
Meaning |
|
Enthalpy of Reaction |
Enthalpy change during a chemical reaction. |
|
Enthalpy of Formation |
Enthalpy change when one mole of a compound forms from its constituent elements in their standard states. |
|
Enthalpy of Vaporisation |
Enthalpy change when a liquid changes into a gas. |
|
Enthalpy of Sublimation |
Enthalpy change when a solid changes directly into a gas. |
|
Enthalpy of Combustion |
Enthalpy change when one mole of a substance burns completely in oxygen. |
Important note: The standard enthalpy of formation of an element in its standard state is zero.
Calculating Enthalpy Changes
|
Method |
Formula |
How to Apply |
|
Using Standard Enthalpies of Formation |
\(\displaystyle \Delta H^\circ_{\text{reaction}}=\sum\Delta H_f^\circ(\text{products})-\sum\Delta H_f^\circ(\text{reactants})\) |
Multiply each formation enthalpy by its stoichiometric coefficient. Subtract the total for reactants from the total for products. |
|
Using Bond Enthalpies |
\(\displaystyle \Delta H\approx\sum E(\text{bonds broken})-\sum E(\text{bonds formed})\) |
Add the energy required to break bonds and subtract the energy released when new bonds form. |
Remember: Breaking bonds requires energy, while forming bonds releases energy. Balance the chemical equation and count the bonds carefully before calculating the enthalpy change.
Spontaneity, Entropy and Gibbs Free Energy
A spontaneous process can proceed naturally under the specified conditions without continuous external intervention. Examples include heat flowing from a hotter body to a colder body and water flowing downhill.
Entropy
-
Definition: Entropy (S) measures the dispersal of energy and the number of accessible microscopic arrangements in a system.
-
Processes that may increase entropy: Evaporation, sublimation and gas expansion often increase entropy under suitable conditions.
-
Entropy change of the universe: The total entropy change is the sum of the entropy changes of the system and its surroundings.
\[ \boxed{\Delta S_{\text{universe}}=\Delta S_{\text{system}}+\Delta S_{\text{surroundings}}} \]
- Condition for spontaneity: A spontaneous process has a positive total entropy change of the universe.
\[ \boxed{\Delta S_{\text{universe}}>0} \]
- Important point: The entropy of the system itself can increase or decrease. Spontaneity depends on the total entropy change of the universe.
Gibbs Free Energy
-
Definition: Gibbs free energy (G) relates enthalpy, entropy and temperature to help determine the thermodynamic direction of a process.
-
Gibbs free energy equation:
\[ \boxed{\Delta G=\Delta H-T\Delta S} \]
- Conditions at constant temperature and pressure:
-
G<0: The process is spontaneous in the forward direction.
-
G=0: The system is at equilibrium.
-
G>0: The forward process is non-spontaneous; the reverse direction is favoured.
- Numerical calculation tip: Keep the units of H and TS consistent before substituting values into the equation. For example, if H is in kJ/mol, convert TS to kJ/mol as well.
Thermodynamics becomes easier to revise when you connect its equations with the physical meaning of heat, work and energy changes. Focus on the first law, pressure-volume work, enthalpy calculations, entropy and Gibbs free energy. Practise numerical problems using the correct sign conventions and units to strengthen your understanding of the chapter for NEET 2027.
Related Resources
Thermodynamics NEET FAQs
What is the first law of thermodynamics?
What is the sign of work during expansion?
Why is internal energy change zero during an isothermal process involving an ideal gas?
What is the difference between reversible and irreversible processes?
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