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Thermodynamics is an important branch of Chemistry that deals with the study of energy transformations and the laws that govern them. You may already know that energy can neither be created nor destroyed—but where does this energy come from? This unit on thermodynamics will answer that question along with many other related concepts. We have provided detailed notes on Class 11 Chemistry (Part I), Chapter 5, to help you understand the fundamental principles of thermodynamics in a clear and structured manner.
Contents
- 1 Introduction
- 2 Thermodynamic Terms
- 3 Applications
- 4 Measurement of ΔU and ΔH: Calorimetry
- 5 Enthalpy Change, ΔrH of a Reaction – Reaction Enthalpy
- 6 Enthalpies for Different Types of Reactions
- 7 Spontaneity
- 8 Entropy and Second & Third Laws of Thermodynamics
- 9 Gibbs Energy Change and Equilibrium
- 10 Important Definitions in NCERT Notes Class 11 Chemistry (Part-I) Chapter 5: Thermodynamics
- 11 FAQs
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Introduction
Different forms of energy (chemical, thermal, mechanical, electrical) transform into one another under various conditions. For example, chemical energy stored in molecules can be released as heat during the combustion of fuels (e.g., methane, LPG, coal). The same chemical energy may be converted to mechanical work (e.g., in engines) or electrical energy (e.g., in dry cells).
Thermodynamics focuses on macroscopic systems containing large numbers of molecules, not on the microscopic mechanisms or rates of transformation. It considers only the initial and final equilibrium states of a system. Laws of thermodynamics describe energy changes only under such equilibrium or quasi‐equilibrium conditions.
Thermodynamic Terms
In this section, we have provided the thermodynamic terms and their definitions.
The System and the Surroundings
In thermodynamics, the system refers to the specific part of the universe under observation. The surroundings include everything else apart from the system. Together, the system and the surroundings constitute the universe:
Universe = System + Surroundings
In practical terms, the surroundings are considered to be the portion of the universe that can interact with the system. Usually, the immediate region around the system forms the surroundings.
Example: If a reaction occurs between substances A and B in a beaker, the beaker and its contents form the system, while the room in which the beaker is placed is the surroundings.
The boundary is the separation between the system and the surroundings. It can be real or imaginary and is used to track the movement of matter and energy.
Types of the System
Systems are classified based on the exchange of energy and matter:
- Open System: Both energy and matter can be exchanged with the surroundings.
Example: Reactants in an open beaker. - Closed System: Energy can be exchanged, but matter cannot.
Example: Reactants in a closed metallic vessel. - Isolated System: Neither energy nor matter is exchanged with the surroundings.
Example: Reactants in a thermos flask or insulated container.
The State of the System
A thermodynamic system is described by specifying measurable properties such as pressure (p), volume (V), temperature (T), and composition. These properties are called state variables or state functions because they depend only on the current state of the system, not on how it reached that state. Not all properties need to be specified to define the system. A minimum set of independent properties can define the state, and all other properties are then fixed.
The Internal Energy as a State Function
The internal energy (U) of a system is the total energy, including chemical, electrical, mechanical, and other forms of energy. It changes when:
- Heat is absorbed or released
- Work is done on or by the system
- Matter enters or leaves the system
(a) Work
In an adiabatic process (no heat exchange), only work changes the internal energy.
If a system changes from state A (temperature TA, internal energy UA) to state B (TB > TA, internal energy UB) by doing work, the internal energy change is:
ΔU = UB − UA = w (for adiabatic process)
Sign convention:
- w > 0: Work is done on the system.
- w < 0: Work is done by the system.
(b) Heat
Heat (q) is the energy transferred between the system and surroundings due to temperature difference.
If only heat is involved (no work), the internal energy change is:
ΔU = q
Sign convention:
- q > 0: Heat is absorbed by the system.
- q < 0: Heat is released by the system.
The General Case
When both heat and work are involved, the internal energy change is given by the First Law of Thermodynamics:
ΔU = q + w
For an isolated system, q = 0 and w = 0, so:
ΔU = 0
Applications
Chemical reactions often involve heat and mechanical work. Thermodynamics helps relate energy changes to internal energy.
Work
When a gas expands or contracts, the work done is:
w = – p_ext × ΔV
(ΔV = Vf − Vi)
- If ΔV > 0 (expansion), then w < 0
- If ΔV < 0 (compression), then w > 0
Reversible Work (external pressure ≈ internal pressure):
w = – ∫ p dV
For isothermal reversible expansion of an ideal gas:
w = – nRT ln(Vf / Vi)
Free Expansion (into vacuum):
p_ext = 0, so w = 0
At constant volume:
w = 0, so ΔU = qV
Enthalpy, H
Enthalpy (H) is defined as:
H = U + pV
At constant pressure, the enthalpy change is:
ΔH = ΔU + pΔV
Since qp = ΔH, enthalpy equals the heat absorbed at constant pressure.
For reactions involving gases:
ΔH = ΔU + ΔngRT, where Δng is the change in moles of gas.
Heat Capacity
Heat capacity (C) is defined as:
q = C × ΔT
Specific heat (c):
q = m × c × ΔT
Molar heat capacity (Cm): heat required for 1 mole of substance.
Cp and Cv for Ideal Gases
At constant volume:
qV = nCvΔT
At constant pressure:
qp = nCpΔT
Relation:
Cp – Cv = R
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Measurement of ΔU and ΔH: Calorimetry
Carried out in a bomb calorimeter, which is a strong steel container (the “bomb”) immersed in water. The substance is ignited electrically. The rise in water temperature helps calculate the heat released:
ΔU Measurements
Done using constant pressure calorimetry (e.g., in a coffee cup). The change in temperature gives the heat exchanged at constant pressure:
qV = ΔU
q = Ccalorimeter × ΔT
ΔH Measurements
Measured using a constant pressure calorimeter (e.g., coffee cup).
qp = ΔH
- Exothermic: ΔH < 0
- Endothermic: ΔH > 0
Enthalpy Change, ΔrH of a Reaction – Reaction Enthalpy
For a reaction:
∑ bi Reactants → ∑ ai Products
The reaction enthalpy is:
ΔrH = ∑ ai H(products) − ∑ bi H(reactants)
Standard Enthalpy of Reactions
ΔrH°: All substances in standard states (1 bar, 298 K).
Enthalpy Changes during Phase Transformations
- Fusion: ΔfusH° > 0
- Vaporisation: ΔvapH° > 0
- Sublimation: ΔsubH° > 0
Standard Enthalpy of Formation
ΔfH° is the enthalpy change for forming 1 mole of a compound from elements.
ΔfH° = 0 for elements in the standard state.
ΔrH° = Σ ΔfH° (products) − Σ ΔfH° (reactants)
Thermochemical Equations
Thermochemical equations include physical states and ΔrH°.
Hess’s Law of Constant Heat Summation
Total enthalpy change is the same, whether the reaction occurs in one or multiple steps.
ΔH(overall) = ΔH1 + ΔH2 + …
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Enthalpies for Different Types of Reactions
The enthalpies for different types of reactions are discussed below.
Standard Enthalpy of Combustion
ΔcH°: Enthalpy change when 1 mole of substance burns completely in oxygen.
Enthalpy of Atomization
ΔaH°: Enthalpy change to convert 1 mole of element into gaseous atoms.
Example: H₂(g) → 2H(g)
Bond Enthalpy
Energy to break 1 mole of specific bonds in gas phase.
ΔrH° ≈ Σ (bond enthalpies of reactants) – Σ (bond enthalpies of products)
Lattice Enthalpy
Energy required to convert 1 mole of ionic solid into gaseous ions.
Enthalpy of Solution
ΔsolH°: Enthalpy change when 1 mole of solute dissolves in a solvent.
Spontaneity
A spontaneous process occurs without external aid. Examples: flow of heat from hot to cold body, diffusion of gas, melting of ice at room temperature. Thermodynamics helps to determine if a process is spontaneous. Spontaneity does not depend solely on energy change (ΔU or ΔH); some endothermic reactions are spontaneous.
Entropy and Second & Third Laws of Thermodynamics
Entropy is a measure of disorder or randomness. It is a state function. Entropy increases with:
- Temperature
- Phase transition: solid → liquid → gas
- Mixing of substances
Third Law of Thermodynamics
The entropy of a perfect crystal at absolute zero (0 K) is zero.
Gibbs Energy Change and Equilibrium
G = H – TS (Gibbs free energy)
ΔG = ΔH – TΔS
- ΔG < 0: Spontaneous
- ΔG > 0: Non-spontaneous
- ΔG = 0: Equilibrium
ΔrG° = – RT ln K, where K is the equilibrium constant.
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Important Definitions in NCERT Notes Class 11 Chemistry (Part-I) Chapter 5: Thermodynamics
Here are the important definitions from NCERT Class 11 Chemistry Chapter 5: Thermodynamics:
- Open System: A system in which both energy and matter can be exchanged with the surroundings. Example: Reactants in an open beaker.
- Closed System: A system in which energy can be exchanged with the surroundings, but matter cannot. Example: A reaction in a closed metallic vessel.
- Isolated System: A system that can exchange neither energy nor matter with the surroundings. Example: A substance in a thermos flask.
- State Function: A property of the system that depends only on its current state and not on how it got there. Examples: Pressure (p), Volume (V), Temperature (T), Internal Energy (U), Enthalpy (H), Entropy (S).
- Adiabatic Process: A process in which no heat is exchanged between the system and its surroundings. Any change in internal energy is due to work only.
- Isothermal Process: A process in which the temperature of the system remains constant during the change.
- Isobaric Process: A process that occurs at constant pressure.
- Isochoric Process: A process that occurs at constant volume.
- Standard State: The most stable form of a substance at 1 bar pressure and a specified temperature (usually 298 K).
- Lattice Enthalpy: The amount of energy required to completely separate one mole of an ionic solid into its gaseous ions.
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FAQs
Answer: Internal energy (U) is the total energy contained within a system, including all forms such as kinetic, potential, chemical, and electrical energy. Enthalpy (H) is defined as the sum of internal energy and the product of pressure and volume of the system:
H = U + pV. Enthalpy is particularly useful for studying heat changes at constant pressure, whereas internal energy is useful for constant volume processes.
Answer: The First Law of Thermodynamics states that energy cannot be created or destroyed, only converted from one form to another. In chemical reactions, this is represented as:
ΔU = q + w, where ΔU is the change in internal energy, q is the heat added to the system, and w is the work done on the system. It ensures that the total energy change in a chemical reaction accounts for both heat exchange and work done.
Answer: The spontaneity of a chemical reaction is determined by the Gibbs free energy change (ΔG), which is given by the equation:
ΔG = ΔH – TΔS, where ΔH is the enthalpy change, T is the temperature, and ΔS is the entropy change.
If ΔG < 0, the reaction is spontaneous.
If ΔG > 0, the reaction is non-spontaneous.
If ΔG = 0, the system is at equilibrium.
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