CHAPTER-5 THERMODYNAMIC

The word thermo means heat and dynamic means flow of motion, these word come form Greek language. The branch of science which deal with different form of energy and their quantitative inter-relationship is called thermodynamic. It is have mainly 4 things, called zeroth, first, second and third laws of thermodynamics. All these laws based on human experience and have no theoretical proof.

GENERAL INTRODUCTION-

Thermodynamic help us to in- predict in extent of chemical reaction, force that derived that chemical reaction, weather related reaction under given temperature and pressure may be or not, etc.

THERMODYNAMIC TERMS-

The part of universe that is under the thermodynamic observation is called system whereas the remaining part of universe which can interact with system called surrounding.

THE UNIVERSE= THE SYSTEM + THE SURROUNDING

The real or imaginary lines which separate two or more system from it surrounding is called boundary.

Types of system-3

Open System- This system can exchange the matter as well as energy with its surrounding and its boundary is not sealed, non-insulated thermally.

Closed System- This system does not exchange matter with its surrounding but can exchange energy from one form to another and its boundary is sealed but not insulated.

Isolated System- This system neither exchange matter nor energy with its surrounding. Here boundary is sealed as well as insulated.

Macroscopic Properties– Thermodynamic deal with matter in bulk and the properties of system arises from collective behaviour of large number of species called macroscopic properties or bulk properties of system.

These are two macroscopic properties-

Intensive Properties- Those property which does not depend on quantity or size of matter. eg- pressure, temperature, density, etc.

Extensive properties- Those property which depend on size or quantity of matter. ex-number of mole, enthalpy, etc.

State of the system-

When the value of macroscopic properties are definite, then the system is called definite state. If the property is change, then the state of system also change. The measurable property is called variable. ex- temperature, pressure, volume etc.

State Functions and path Functions –

When the variable depend on the initial and final state of system, but not depend on path, is called state function or state variable. ex- pressure, volume, enthalpy etc.

The variable affected by a followed path, is called path function. ex- work(W), heat(q), etc.

Types of Thermodynamics Processes-

1. Adiabatic Process– The process which does not exchange heat with its surrounding.
   • In this system have a wall which separate system or surrounding is called adiabatic wall .
2. Isothermal Process-In this process, temperature remain constant or fixed.
3. Isochoric Process– In this process, volume remain constant.
4. Isobaric Process-Change of state is brought about at constant pressure.
5. Cyclic Process-In this process, the system undergoes different changes and finally return to its initial stages.
6. Reversible Process- This process is very slow, allowing the system and surroundings to remain in equilibrium at every step, and be reversed without leaving any change in either.
7. Irreversible Process– It is real process that occur quickly, causing the system and surrounding lose equilibrium and resulting in permanent change that cannot be completely reversed.

INTERNAL ENERGY AS A STATE FUNCTION-

1. Internal Energy (U):
   • Definition: Represents the total energy of a system, which includes chemical, electrical, mechanical, or other types of energy.
   • Change in Internal Energy (∆U): Can occur due to:
     • Heat exchange with the surroundings.
     • Work done on/by the system.
     • Matter entering or leaving the system.
2. Work and Adiabatic Systems:
   • An adiabatic system is one where no heat exchange occurs between the system and its surroundings. The wall separating the system from the surroundings is called the adiabatic wall.
   • The adiabatic process refers to the change in the state of a system where no heat is exchanged.
   • Example: A system with water in an insulated container (adiabatic system) can have its internal energy changed by mechanical work (e.g., rotating paddles or electrical work using an immersion rod).
   • Result of Work: When work is done on the system, the internal energy increases, leading to a temperature rise.
   • Joule’s Experiments: J.P. Joule demonstrated that the same amount of work (irrespective of how it is done) produces the same temperature change in the system, showing that internal energy change is independent of the path taken.
3. Internal Energy as a State Function:
   • State Function: A property whose value depends only on the current state of the system and not on how the system reached that state.
   • Equation for Internal Energy Change: The change in internal energy due to work (in an adiabatic process) is given by:
     • ΔU=U2−U1=wad​ (where wad is the adiabatic work).
   • Examples of State Functions: Temperature (T), volume (V), pressure (p) are state functions, meaning their changes are independent of the path taken.
4. Heat Transfer and Internal Energy:
   • Heat (q): The transfer of energy due to a temperature difference between the system and its surroundings.
   • Heat Transfer Process: If the system allows heat exchange, the heat absorbed or released by the system changes its internal energy.
   • Thermally Conducting System: In a system with thermally conducting walls, heat is transferred to/from the system and changes its internal energy without doing work.
   • Heat and Internal Energy Change: When heat is transferred from surroundings to the system, the internal energy increases (q > 0). When heat is transferred from the system to surroundings, the internal energy decreases (q < 0).
5. General Case of Energy Change:
   • In a general case where both heat and work are involved in changing the state of the system, the change in internal energy is expressed as:
     • ΔU=q+w
   • First Law of Thermodynamics: This equation represents the first law, which states that energy cannot be created or destroyed, only transformed:
     • Law of Conservation of Energy: The total energy of an isolated system remains constant.
6. Key Differences between Thermodynamic and Mechanical Properties:
   • Volume: The absolute value of volume is easily measurable in a system.
   • Internal Energy: The absolute value of internal energy cannot be measured directly; we can only measure the change in internal energy (∆U).

Some Appications-

Work Done In Thermodynamics-

• Mechanical Work: This refers to pressure-volume work (i.e., work done by the gas as it expands or contracts).
• Compression: If an ideal gas is compressed by an external pressure greater than the internal pressure, the work done on the system is calculated using this formula
  • W = -P_ext ΔV= -P_ext (V_f – V_i ), where P_ext is the external pressure, and (V_f – V_i) are the final and initial volumes, respectively.
• Reversible Processes: These occur when the system is in equilibrium with the surroundings at every point. The work done during reversible compression or expansion can be expressed using the pressure and volume-
  • W_rev = – _vi∫^vf  P_ext​dV
  • Free Expansion: In free expansion, the gas expands into a vacuum (no external pressure), so no work is done (W = 0).
• 2. Isothermal and Adiabatic Processes–
  • Isothermal (Constant Temperature)- For an isothermal expansion of an ideal gas, the internal energy change (ΔU) is zero because temperature is constant. The heat absorbed by the gas is equal to the work done-
    • Q= -W
    • For a reversible isothermal expansion-
      • W = _nRT In(V_f / V_i)
  • diabatic Process (No Heat Exchange): Here, no heat is exchanged (Q = 0), and the work done changes the internal energy-
    • ΔU = W
• 3. Internal Energy (ΔU) and Enthalpy (ΔH)
  • Internal Energy: In thermodynamics, the internal energy change ΔU\Delta UΔU is the sum of heat and work interactions with the system.
    • If ΔV=0 = 0, then ΔU=Q_v​ (at constant volume).
  • Enthalpy: Enthalpy HHH is a useful state function for reactions at constant pressure. It is defined as-
  • H = U + _pV
    • The change in enthalpy at constant pressure is-
  • ΔH = ΔU + Δ (pV)
  • For reactions involving gases, this is useful for calculating the heat exchanged.
• Relation between ΔH and ΔU- For reactions involving gases, we can express the difference in enthalpy and internal energy as-
  • ΔH = ΔU + Δn_g​RT
  • where Δn_g​ is the change in the number of moles of gas.
• 4. Heat Capacity
• Heat Capacity (C): This is the amount of heat required to raise the temperature of a substance by 1°C (or 1 K). The relationship is-
  • Q = C ΔT
• where C is the heat capacity and ΔT is the temperature change.
• Molar Heat Capacity: For one mole of a substance, the heat capacity is called the molar heat capacity C_m​, and it is the heat required to raise the temperature of one mole by 1°C.
• Specific Heat: Specific heat is the heat capacity per unit mass of the substance. It is the heat required to raise the temperature of 1 gram of the substance by 1°C
• 5. Relationship between C_p and C_v for Ideal Gases-
  • C_p is the heat capacity at constant pressure, and C_v​ is the heat capacity at constant volume. The relationship between them for an ideal gas is-
    • ​C_p + C_v = R, where R is the universal gas constant.
• 6. Applications in Competitive Exams
• Isothermal and Adiabatic Processes: Problems in exams often ask you to calculate work done, heat absorbed, or internal energy change in these processes. Be familiar with the equations and the underlying concepts.
• Calculating Heat and Work: Understand how to calculate work done by gases, especially in processes where volume changes under constant external pressure, or when the process is reversible.
• Ideal Gas Law: Often, the equations involving pV = nRT will be central in solving problems about work and heat. Make sure to practice using the ideal gas law to derive other thermodynamic relations.
• Free Expansion (No Work Done): If a gas expands in a vacuum, there’s no work because Pext = 0, and in isothermal expansion, the heat absorbed equals the work done.
• Reversible Isothermal Expansion: For a reversible isothermal process, use the formula for work, and remember that the internal energy change is zero.
• Final Tips:
• Focus on the basic definitions and equations for each process (isothermal, adiabatic, reversible, irreversible).
• Practice using equations like W =−Pext ΔV and the relationship between ΔH and ΔU in various conditions.

Measurement of ∆U and Calorimetry (Simplified)

Calorimetry: Energy Measurement Technique

• Definition: Calorimetry is a technique to measure energy changes (heat) in chemical or physical processes.
• Setup:
  • Uses a calorimeter, immersed in a known amount of liquid.
  • By knowing the heat capacities of the liquid and calorimeter, and monitoring temperature changes, heat transfer in the process can be calculated.
• Conditions:
  • Constant Volume (q_v): No volume change; often used in sealed setups like bomb calorimeters.
  • Constant Pressure (q_p​): Pressure remains constant; commonly used in open systems under atmospheric pressure.

(a) ∆U (Internal Energy Change) Measurements

• Bomb Calorimeter:
  • A steel vessel (“bomb”) is used, immersed in a water bath to prevent heat loss.
  • Combustible substances burn in pure oxygen inside the bomb.
  • Heat released is transferred to the surrounding water, and the temperature rise is recorded.
  • Since the volume is constant, no work is done (W=0).
  • Temperature change is converted into heat (q_v​) using the calorimeter’s heat capacity.
  • Key Equation: q_v=C⋅ΔT, where C= heat capacity, ΔT= temperature change.

(b) ∆H (Enthalpy Change) Measurements

• Constant Pressure Setup:
  • Measures heat (q_P​) at constant pressure, often under atmospheric conditions.
  • Enthalpy change (ΔH) = Heat of reaction (q_P​).
• Heat Significance:
  • Exothermic reactions: Heat released, ΔH<0.
  • Endothermic reactions: Heat absorbed, ΔH>0.

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Reaction Enthalpy (ΔrH)

• Definition: Change in enthalpy when reactants turn into products.
• Formula: ΔrH=Sum of enthalpies of products−Sum of enthalpies of reactants
• Applications:
  • Used to calculate energy needs for maintaining industrial reactions.
  • Determines temperature dependence of equilibrium constants.

Standard Enthalpy (ΔH^∘)

• Standard Conditions:
  • Substances are in their standard state: pure form at 1 bar and specified temperature (commonly 298 K).
• Types:
  • Fusion (Δ_fusH^∘): Energy to melt 1 mole of solid at standard state.
  • Vaporization (Δ_vapH^∘): Energy to vaporize 1 mole of liquid.
  • Sublimation (Δ_subH^∘): Energy for 1 mole of solid to directly turn into vapor.

Key Equations

1. Energy Change (ΔrH^∘):
   • Δ_rH^∘= ∑Δ_fH^∘(product)- ∑Δ_fH^∘(reactant)-
   • Δf​H^∘: Standard molar enthalpy of formation.
2. Internal Energy and Enthalpy Relation:
   • ΔU = ΔH − _pΔV, where-
     • ΔU: Internal energy change
     • ΔH: Enthalpy change.
     • _pΔV: Work done due to volume change.

Phase Changes

• Melting (Fusion):
  • Example: Ice to water at 273 K.
  • Δ_fus​H^∘ > 0 (endothermic).
• Evaporation (Vaporization):
  • Example: Water to steam
  • Δ_vap​H^∘ > 0.
• Sublimation:
  • Example: Dry ice to gas.
  • Δ_sub​H^∘ > 0.

Practical Problem

1. Bomb Calorimeter Example:
   • 1 g of graphite burns to produce CO₂ at constant volume.
   • Heat (q_v) = −20.7 kJ (exothermic, heat released).
   • Combustion of 1 mole graphite: ΔU=−248 kJ/mol.
2. Water Evaporation:
   • To evaporate 18 g water at 298 K:
     • Heat required=n⋅Δ_vapH^∘=1 * 44.01  kJ / mol = 44.01 kJ.

Thermochemical Equations and Related Concepts

Thermochemical Equations

1. A thermochemical equation includes the balanced chemical equation and the enthalpy change (ΔrH).
   • Example: C₂H₅OH_(l)+3O₂(g)→2CO₂(g)+3H₂O_(l), ΔrH^∘ =−1367 
   • A negative ΔrH implies the reaction is exothermic.

2. Key Conventions:

• Coefficients in the equation represent moles, not molecules.
• ΔrH∘ refers to the enthalpy change for the specific moles indicated and is expressed in kJ/mol.
• Enthalpy is extensive, meaning its value depends on the number of moles.

3. Reversing a chemical equation changes the sign of ΔrH∘.

Example:
N₂(g)+3H₂(g)→2NH₃(g);ΔrH^∘=−91.8 kJ / mol
Reversed: 2NH₃(g)→N₂(g)+3H₂(g);ΔrH^∘=+91.8 kJ / mol

Hess’s Law of Constant Heat Summation

• The total enthalpy change for a reaction is the sum of enthalpy changes for its intermediate steps.
• This is possible because enthalpy is a state function, depending only on initial and final states, not the path.

Application Example:
To find ΔrH∘\Delta_rH^\circΔr​H∘ for a reaction indirectly:

1. Use related reactions with known ΔrH^∘.
2. Reverse or scale these reactions as needed.
3. Sum the adjusted enthalpy changes to get the desired ΔrH^∘.

Types of Enthalpies

1. Standard Enthalpy of Combustion (ΔcH∘\Delta_cH^\circΔc​H∘)
   • The heat released when 1 mole of a substance undergoes complete combustion under standard conditions.
   • Example: C₄H₁₀+6.5O₂→4CO₂+5H₂O, ΔcH^∘=−2658.0 kJ/mol.
2. Enthalpy of Atomization (ΔaH∘\Delta_aH^\circΔa​H∘)
   • Energy required to break all bonds in 1 mole of a compound to form gaseous atoms.
   • Example: CH_4(g) → C_(g)+ 4H_(g), ΔaH^∘ =1665 kJ/mol.
   • For diatomic molecules, this is the same as bond dissociation enthalpy.
3. Bond Enthalpy (Δ_bondH^∘)
   • Energy required to break 1 mole of a specific bond in a gaseous compound.
   • Mean bond enthalpy: Average energy to break a particular bond type in a polyatomic molecule.
   • Example: Mean Δ _C–HH∘ in CH₄ = 416  kJ/mol.
4. Lattice Enthalpy
   • The energy required to separate 1 mole of an ionic compound into its gaseous ions.
   • Example: NaCl_(s)→Na⁺_(g)+Cl⁻_(g), Δ_lattice​H^∘=+788kJ/mol.

Bond Dissociation vs. Formation:

• Energy is absorbed to break bonds (endothermic).
• Energy is released when bonds form (exothermic).

Using Hess’s Law and Bond Enthalpies:

• Δr​H^∘= ∑Bond Energies of Reactants − ∑Bond Energies of Products.
• Useful for reactions where direct enthalpy values are unavailable.

Practical Applications:

• Combustion for energy (e.g., fuels, cooking gas).
• Understanding biological energy transformations.

Spontaneity and Thermodynamics

Spontaneous Processes

1. Definition: A process is spontaneous if it can occur without external help. However, spontaneity doesn’t indicate speed (e.g., hydrogen and oxygen form water slowly but spontaneously).
2. Unidirectionality: Spontaneous changes naturally proceed in one direction (e.g., gas expansion, heat flow from hot to cold). Reverse changes require external intervention.
3. Key Criterion: Spontaneous processes are irreversible without external factors.

Energy and Spontaneity

1. Enthalpy (ΔH\Delta HΔH):
   • Exothermic reactions (ΔH < 0) often indicate spontaneity (e.g., water formation releases heat).
   • Endothermic reactions (ΔH > 0) can also be spontaneous (e.g., nitrogen and oxygen forming NO₂).
   • Hence, a decrease in enthalpy supports spontaneity but isn’t the sole factor.
2. Entropy (ΔS):
   • Entropy measures randomness or disorder in a system.
   • Spontaneous changes increase total entropy in an isolated system (Second Law of Thermodynamics).
   • Higher disorder (e.g., gases diffusing or solids melting) corresponds to higher entropy.
   • Entropy change depends on heat transfer and temperature (ΔS = q_rev / T).
     

Entropy in Reactions

1. Qualitative Understanding:
   • Gases have higher entropy than liquids or solids.
   • Breaking larger molecules into smaller ones or forming gases increases entropy.
2. Quantitative Understanding:
   • ΔS_total ​= ΔS_system​+ΔS_surroundings​
   • If ΔS_total>0, the process is spontaneous.
3. Surroundings Contribution:
   • The entropy change in surroundings depends on the system’s heat loss or gain:
   • ΔS_surroundings ​= −ΔH_system​/T​

1. Gibbs Free Energy (ΔG)
   • Definition: Combines enthalpy and entropy to determine spontaneity:
     • ΔG = ΔH − T ΔS
   • Represents the usable energy for work.
2. Spontaneity Conditions:
   • ΔG<0: Spontaneous process.
   • ΔG>0: Non-spontaneous process.
   • ΔG=0: System is in equilibrium
3. Temperature Dependency:
   • Spontaneity depends on the balance between ΔH and TΔS:
   • Exothermic and increasing entropy (ΔH<0, ΔS>0): Always spontaneous.
   • Endothermic but highly increasing entropy (ΔH>0, ΔS>0): Spontaneous at high temperatures.

Second Law of Thermodynamics

• Statement: In an isolated system, the total entropy always increases for a spontaneous process.
• Implications:
  • Exothermic reactions often drive spontaneity by increasing surroundings’ entropy.
  • The natural direction of change aligns with increased disorder.

Key Takeaways for Competitive Exams

1. Enthalpy and entropy are not independently sufficient to predict spontaneity; use ΔG=ΔH−TΔS.
2. Understand entropy trends: gases > liquids > solids.
3. Analyze temperature effects on reaction feasibility.
4. Learn entropy-related concepts like disorder, randomness, and heat distribution.
5. Use Gibbs energy as the ultimate criterion for spontaneous reactions in practical systems.

Absolute Entropy and Third Law of Thermodynamics

1. Molecular Motions and Temperature Effects:
   • Molecules exhibit three types of motions:
     • Translational: Straight-line movement.
     • Rotational: Spinning motion.
     • Vibrational: Stretching and compressing of molecular bonds.
   • Higher temperature increases the intensity of these motions, leading to greater entropy (disorder).
   • Lower temperature reduces these motions, decreasing entropy.
2. Third Law of Thermodynamics:
   • The entropy of a pure crystalline substance approaches zero as the temperature approaches absolute zero (0 K).
   • This happens because, at absolute zero, the crystal structure has perfect order, resulting in zero entropy.
   • Limitation: This law is valid for pure crystalline solids. Solutions and supercooled liquids do not have zero entropy at 0 K due to structural imperfections.
3. Importance of the Third Law:
   • Allows calculation of the absolute entropy of pure substances using thermal data from 0 K to a specific temperature (e.g., 298 K).
   • Standard entropies are useful for determining standard entropy changes in reactions via calculations similar to Hess’s law.

Gibbs Free Energy, Equilibrium, and Spontaneity

1. Significance of Gibbs Energy (ΔG):
   • Spontaneity: A reaction is spontaneous if ΔG < 0.
   • Equilibrium: At equilibrium, ΔG = 0.
   • Helps predict the maximum useful work obtainable from a reaction.
2. Reversible Reactions and Dynamic Equilibrium:
   • In a reversible reaction, the system remains in equilibrium with its surroundings.
   • At equilibrium, the system’s free energy is at its minimum. Any deviation would spontaneously move back to this minimum energy state.
3. Gibbs Energy and Equilibrium Constant (K):
   • Relationship: Relationship: ΔrG⁰=−RTlnK or ΔrG⁰= -2.3 03 RT log K
   • For endothermic reactions (positive ΔH°), K≪1, indicating minimal product formation.
   • For exothermic reactions (negative ΔH°), K≫1, indicating significant product formation.
   • Entropy changes (ΔS°) also influence K, affecting the reaction’s extent based on whether ΔS° is positive or negative.
4. Temperature Effects:
   • Reactions may become spontaneous or non-spontaneous depending on the interplay of ΔH° and ΔS° at different temperatures:
   • High temperatures favor reactions with positive ΔS°.
   • Low temperatures favor reactions with negative ΔS°.
     

Calculations and Practical Applications:

1. Determining ΔG° and K:
   • Gibbs free energy change can be calculated using enthalpy (ΔH°) and entropy (ΔS°) data.
   • ΔG⁰=ΔH⁰−TΔS⁰
   • Equilibrium constant K is derived from ΔG° at specific temperatures for optimizing reaction conditions.
2. Use Cases:
   • Knowing K, ΔG° can be calculated for different temperatures to predict reaction feasibility.
   • Standard entropy changes allow us to estimate reaction spontaneity and equilibrium behavior.
3. Example Interpretation:
   • Strongly exothermic reactions typically have large KKK, meaning they proceed nearly to completion.
   • Endothermic reactions may require specific conditions (e.g., high temperature) to proceed significantly.

Summary of Key Thermodynamic Principles:

1. Thermodynamics: Studies energy changes in chemical and physical processes.
   • Divides the universe into a system and its surroundings.
   • Heat (q) and work (w) relate to internal energy (ΔU) via ΔU = q + w.
2. Key Functions:
   • Enthalpy (H): Heat content at constant pressure.
   • Entropy (S): Measure of disorder or randomness.
   • Gibbs Free Energy (G): Combines enthalpy and entropy to determine spontaneity.
3. Reaction Spontaneity:
   • Spontaneous reactions increase the total entropy of the universe.
   • Gibbs energy provides a practical measure to predict spontaneity at constant pressure.

By understanding these principles and their interconnections, you can solve competitive thermodynamics questions effectively.

These all are the notes of chapter 5 chemistry. And BELOW HERE have important questions which help in understand whole chapter. *#THANKS FOR VISITING, VISIT AGAIN#* 😊

IMPOTANT QUESTIONS AND ANSWER ARE-

Questions And Answer In MCQs Form-

1. The branch of science that deals with energy and its interconversions is known as:
   A) Thermometry
   B) Thermodynamics
   C) Thermochemistry
   D) Thermophysics
   Answer: B

2. Which of the following is NOT a law of thermodynamics?
A) Zeroth law
B) First law
C) Third law
D) Law of motion
Answer: D

3. In thermodynamics, the part of the universe under study is called the:
   A) System
   B) Surrounding
   C) Boundary
   D) Universe
   Answer: A
4. A system that exchanges neither matter nor energy with its surroundings is called:
   A) Open system
   B) Closed system
   C) Isolated system
   D) Thermodynamic system
   Answer: C
5. Which of the following is an example of an open system?
   A) A sealed bottle of water
   B) A boiling pot of water without a lid
   C) A thermos flask
   D) A bomb calorimeter
   Answer: B

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Processes

6. In which thermodynamic process does temperature remain constant?
   A) Isothermal
   B) Adiabatic
   C) Isochoric
   D) Isobaric
   Answer: A
7. In an adiabatic process:
   A) Heat is exchanged with the surroundings
   B) Volume remains constant
   C) No heat is exchanged with the surroundings
   D) Pressure remains constant
   Answer: C
8. Which process occurs when the system undergoes different changes and finally returns to its initial state?
   A) Adiabatic
   B) Isobaric
   C) Cyclic
   D) Reversible
   Answer: C

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Properties and Functions

9. Which of the following is an intensive property?
   A) Volume
   B) Temperature
   C) Enthalpy
   D) Number of moles
   Answer: B
10. A property that depends only on the initial and final states, not on the path, is called a:
    A) Path function
    B) State function
    C) Macroscopic property
    D) Bulk property
    Answer: B
11. Which of the following is an extensive property?
    A) Pressure
    B) Density
    C) Volume
    D) Temperature
    Answer: C

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Laws and Equations

12. The first law of thermodynamics is expressed as:
    A) ΔU = q + w
    B) ΔU = q − w
    C) ΔU = w − q
    D) ΔU = q × w
    Answer: A
13. For an isothermal process in an ideal gas, which of the following is true?
    A) ΔU = 0
    B) W = 0
    C) q = 0
    D) pV = constant
    Answer: A
14. Which thermodynamic property cannot be directly measured?
    A) Pressure
    B) Volume
    C) Temperature
    D) Internal energy
    Answer: D

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Heat, Work, and Energy

15. In an isothermal reversible expansion of an ideal gas, the work done is given by:
    A) W = nRT ln(Vf/Vi)
    B) W = −PextΔV
    C) W = ΔU − q
    D) W = q − ΔU
    Answer: A
16. During free expansion of a gas:
    A) Work done is maximum
    B) Work done is zero
    C) Heat is absorbed
    D) Heat is released
    Answer: B

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Entropy and Spontaneity

17. Entropy is a measure of:
    A) Heat content
    B) Randomness or disorder
    C) Energy conservation
    D) Work done
    Answer: B
18. For a spontaneous process, the Gibbs free energy change (ΔG) is:
    A) Positive
    B) Zero
    C) Negative
    D) Undefined
    Answer: C

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Applications

19. In a bomb calorimeter, which thermodynamic property is directly measured?
    A) Enthalpy change
    B) Internal energy change
    C) Work done
    D) Heat capacity
    Answer: B
20. The relationship between Cp and Cv for an ideal gas is:
    A) Cp − Cv = R
    B) Cp + Cv = R
    C) Cp/Cv = R
    D) CpCv = R
    Answer: A

Questions And Answers-

1. What does “thermodynamics” mean?

Answer:
The word “thermo” means heat, and “dynamic” means motion or flow. Thermodynamics is the science that studies different forms of energy and their relationships.

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2. What are the main laws of thermodynamics?

Answer:
There are four laws:

• Zeroth Law: Deals with temperature equilibrium.
• First Law: Energy cannot be created or destroyed (conservation of energy).
• Second Law: Entropy (disorder) always increases in an isolated system.
• Third Law: As temperature approaches absolute zero, entropy approaches a minimum value.

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3. What is a system in thermodynamics?

Answer:
A system is the part of the universe we study in thermodynamics. Everything outside the system is called the surroundings.

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4. What are the types of systems?

Answer:

• Open System: Exchanges both matter and energy with surroundings (e.g., boiling water).
• Closed System: Exchanges only energy, not matter (e.g., a sealed, heated container).
• Isolated System: Neither exchanges matter nor energy (e.g., a thermos flask).

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5. What is the difference between intensive and extensive properties?

Answer:

• Intensive Properties: Do not depend on the system size (e.g., temperature, pressure).
• Extensive Properties: Depend on the system size (e.g., volume, enthalpy).

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6. What is the difference between state and path functions?

Answer:

• State Function: Depends only on the system’s initial and final states (e.g., temperature, pressure).
• Path Function: Depends on the path taken to reach the state (e.g., work, heat).

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7. What are some key thermodynamic processes?

Answer:

• Adiabatic: No heat exchange.
• Isothermal: Temperature remains constant.
• Isochoric: Volume remains constant.
• Isobaric: Pressure remains constant.
• Cyclic: System returns to its initial state.

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8. What is internal energy?

Answer:
Internal energy is the total energy of a system, including chemical, mechanical, and other forms. It can change by heat transfer or work done on/by the system.

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9. What does the first law of thermodynamics state?

Answer:
The energy of a system changes due to heat (q) and work (w):
ΔU=q+w
Where ΔU is the change in internal energy.

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10. What is entropy?

Answer:
Entropy measures the disorder or randomness of a system. Spontaneous processes increase the total entropy in an isolated system.

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11. What is Gibbs Free Energy (G)?

Answer:
Gibbs Free Energy determines whether a process is spontaneous:
ΔG= ΔH−TΔS

• ΔG<0: Spontaneous process.
• ΔG>0: Non-spontaneous process.
• ΔG=0: System is in equilibrium.

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12. What is heat capacity?

Answer:
Heat capacity (C) is the heat needed to raise the temperature of a substance by 1°C.

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13. What is the second law of thermodynamics?

Answer:
Entropy of an isolated system always increases in a spontaneous process.

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14. What is the third law of thermodynamics?

Answer:
As temperature approaches absolute zero, a system’s entropy approaches its minimum value.

Questions In Numerical Form-

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