Thermodynamics is a chapter in the CBSE Class 11 Chemistry syllabus from Chemistry Part - I. This chapter hub brings together revision notes, practice questions, worksheets, flashcards, formula sheet to help students learn, practice, and revise Thermodynamics effectively.

Scroll down to find Thermodynamics notes, practice questions, worksheets, and revision resources — all in one place. Use the sidebar to jump to any section, or browse the full page below.

Thermodynamics

NCERT Class 11 Chemistry Chapter 5: Thermodynamics (Pages 136–167)

Class 11 CBSE hub Chemistry chapters Chemistry Part - I

Listen to this article

Summary of Thermodynamics

00:00

00:0000:00

-- min

Thermodynamics at a Glance

Board

CBSE

Class

Class 11

Subject

Chemistry

Book

Chemistry Part - I

Chapter

Pages

136–167

Resources

7 study resources

Thermodynamics Summary

Thermodynamics is a fundamental concept in chemistry that focuses on the energy transformations during chemical reactions. It is essential for students to grasp the concepts of systems and surroundings, which divide the universe into interacting parts. A system could be open, closed, or isolated depending on interactions with the surroundings. Open systems allow exchange of both energy and matter, while closed systems allow only energy exchange and isolated systems allow neither. One crucial concept covered in this chapter is internal energy, which totalizes all forms of energy within a system and changes according to heat and work interactions. The first law of thermodynamics, which states that energy can neither be created nor destroyed, forms the basis for understanding these energy changes. Additionally, we explore state functions, which depend only on the system's current state rather than how that state was achieved. Further, the chapter delves into internal energy changes, exploring how work and heat contribute to these changes. It emphasizes how thermodynamic principles apply primarily during equilibrium states where macroscopic properties such as temperature and pressure remain constant. We will also study enthalpy, which is a convenient way to express energy changes at constant pressure, and explains the importance of heat of reactions, particularly during phase changes and under different conditions. Enthalpy changes can be calculated using various methods such as Hess's law, allowing us to predict reaction behaviors accurately. Additionally, spontaneity and its relation to entropy are introduced. While energy decreases could indicate spontaneity, the role of entropy, which measures disorder or randomness, is more definitive. The second law of thermodynamics emphasizes that processes increase the total entropy of the universe, which serves as a criterion for spontaneity. Lastly, Gibbs free energy combines both enthalpy and entropy to determine the spontaneity of reactions at constant temperature and pressure. Understanding the Gibbs free energy change will aid in predicting the feasibility of reactions based on enthalpy and entropy changes. The relationship between free energy and the equilibrium constant allows chemists to deduce the outcome of reactions at equilibrium. Overall, a thorough understanding of thermodynamics lays the groundwork for students to rationalize chemical phenomena and manipulate them in practical applications.

Thermodynamics Revision Guide

Download the Thermodynamics revision guide with key points, summaries, and quick revision notes for CBSE Class 11 Chemistry.

Key Points

Thermodynamics studies energy transformations.

It analyzes how different forms of energy interact and convert, laying the foundation for chemical processes.

System vs. Surroundings clarification.

A system is the part of the universe being studied; surroundings are everything else, affecting or affected by the system.

Types of systems: Open, Closed, Isolated.

Open systems exchange matter and energy; closed systems exchange only energy; isolated systems exchange nothing.

Internal Energy (U) is a state function.

It represents the total energy within a system, which can change with heat, work, or mass transfer.

First Law of Thermodynamics: ΔU = q + w.

The change in internal energy is equal to heat added to the system plus work done on the system.

Heat (q) can change U.

Heat exchange occurs when there is a temperature difference between a system and its surroundings.

Work (w) is path-dependent.

Work done on/by the system contributes to internal energy changes, but depends on the pathway taken.

Enthalpy (H) is defined: H = U + pV.

It offers a practical measure of heat transfer under constant pressure; ΔH = ΔU + pΔV.

Standard Enthalpy changes defined.

Standard enthalpy values express the heat released or absorbed during reactions at standard conditions.

Hess's Law: ΔH is path-independent.

The total enthalpy change in a reaction is equal to the sum of the enthalpy changes in the individual steps, regardless of the path taken.

Entropy (S) measures disorder.

Entropy quantifies the degree of randomness in a system, with spontaneous processes leading to increased entropy.

Second Law of Thermodynamics.

In isolated systems, total entropy tends to increase, driving spontaneous processes to higher disorder.

Gibbs Free Energy (G): G = H - TS.

It determines spontaneity: if ΔG < 0, the process is feasible; equilibrium occurs when ΔG = 0.

Relationship: ΔG = ΔH - TΔS.

This equation links changes in enthalpy and entropy to the spontaneity of a process.

Equilibrium constant (K) linked to ΔG.

The standard free energy change is related by ΔG° = -RT ln K, where R is the gas constant.

Spontaneous vs. non-spontaneous reactions.

A spontaneous reaction occurs without external intervention; a non-spontaneous one requires input.

Thermodynamic stability indicated by ΔG.

Stable reactions usually have negative ΔG, implying they proceed forward without additional energy.

Entropy is a state function.

Entropy depends only on the initial and final states of the system, not on the pathway taken.

Key thermodynamic terms: ΔU, ΔH, ΔS.

Understand these changes and their signs to predict the behavior of reactions and systems.

Heat capacity related to q = CΔT.

Heat capacity indicates how much heat is needed to change a substance's temperature, essential for calorimetry.

Thermodynamics Practice Questions & Answers

Practice important questions and exam-style problems from Thermodynamics. These questions cover key topics from the CBSE Class 11 Chemistry syllabus.

How to practice: Start with the questions below to test your understanding of Thermodynamics. Use the revision guide to review concepts you find difficult, then come back and retry the questions for better retention.

What is the change in internal energy when no heat is added to the system, but work is done on it?

Single Answer MCQ

Q-00054749

View explanation

If a system does work on its surroundings and no heat is exchanged, what equation describes the change in internal energy?

Single Answer MCQ

Q-00054750

View explanation

When heat is removed from a system and no work is done, what can be said about the change in internal energy?

Single Answer MCQ

Q-00054751

View explanation

What type of wall would a system have if it cannot exchange heat but can do work?

Single Answer MCQ

Q-00054752

View explanation

In a closed system, if a gas expands against a constant external pressure, which of the following statements is TRUE?

Single Answer MCQ

Q-00054753

View explanation

If 100 J of work is done on a system and no heat is transferred, what is the change in internal energy?

Single Answer MCQ

Q-00054754

View explanation

In which situation is the internal energy of a system unchanged?

Single Answer MCQ

Q-00054755

View explanation

If the internal energy of a system decreases due to work done by the system, which of the following must be true?

Single Answer MCQ

Q-00054756

View explanation

View all 100 Thermodynamics questions

In a gas system, if 500 J of heat is added and the system does 200 J of work, what is the change in internal energy?

The mathematical expression ∆U = q + w is associated with which fundamental principle?

What type of system exchanges both energy and matter with its surroundings?

If work is done on the system and heat is lost, how would the internal energy change?

What is the effect on internal energy if a perfectly insulated system expands against an external pressure?

A piston compresses an ideal gas. If the work done on the gas is positive, what happens to its internal energy?

When examining a reaction that produces a gas in a closed container, what concept helps predict the pressure change as the volume changes?

What is a thermodynamic system?

Which of the following describes an open system?

What is the primary purpose of calorimetry in scientific experiments?

In a closed system, what can be exchanged with the surroundings?

In which type of calorimeter is the change in internal energy (∆U) measured?

Which of the following is an example of an isolated system?

How is heat absorbed by a calorimeter calculated?

What does the state of a thermodynamic system depend on?

What does a negative value of ∆H indicate about a reaction?

What is the role of boundaries in a thermodynamic system?

What is the relationship between ∆U and ∆H at constant pressure?

Which statement is true about state functions?

Which factor does NOT affect the heat capacity of a calorimeter?

The heat content of a system at constant pressure is termed as:

At what conditions is the measurement of enthalpy (∆H) commonly performed?

Which of the following is NOT a state variable?

What happens to the heat of reaction (qp) in an endothermic process?

An endothermic reaction is characterized by:

What is the effect of combining two endothermic reactions in a calorimeter?

In thermodynamics, which term describes the total energy of a system?

If the temperature in a bomb calorimeter raises by 1 K and the heat capacity is 15.5 kJ/K, what is the heat released?

How does an isolated system differ from a closed system?

What is calculated when measuring qp in a calorimeter reacting at constant pressure?

Which of the following best describes a state variable in a gas?

Which of the following is NOT a function of calorimetry?

The first law of thermodynamics is essentially a statement of the conservation of:

What is true about the relationship between qv and qp?

If the internal energy of a system increases, what must happen to its surroundings?

In the energy balance of an endothermic reaction, what must occur to maintain constant pressure?

Which term describes the heat absorbed or released during a reaction at constant pressure?

What does the symbol ∆rH represent in a chemical reaction?

In which scenario is the enthalpy change (∆rH) of a reaction positive?

What is the enthalpy change for the reaction: C (s) + O2 (g) → CO2 (g)?

Which factor does not affect the value of enthalpy change, ∆rH?

When calculating ∆rH using Hess's law, what must be true about the final reaction?

What happens to the enthalpy change (∆rH) when a reaction is reversed?

If the enthalpy of the products is higher than that of the reactants, this indicates a:

For the reaction CO (g) + 1/2 O2 (g) → CO2 (g), what is the value of ∆rH?

What does it mean if a reaction has a negative ∆rH and a positive ∆S?

Which formula correctly represents the calculation of enthalpy change?

Combustion of a substance typically results in which sign of ∆rH?

If the enthalpy change for a certain reaction is zero, what can be inferred about the reaction?

What is the correct interpretation of Hess's law in relation to enthalpy?

In an endothermic reaction, if ∆H is greater than the energy changes of the surroundings, what will happen?

What is the standard enthalpy of combustion for one mole of butane?

Which of the following statements best describes the standard enthalpy of formation?

What does a negative value of ∆cH⁰ indicate about a combustion reaction?

If the enthalpy of combustion of glucose is -2802.0 kJ/mol, what does this imply?

Which equation correctly illustrates the relationship between standard enthalpies of different routes leading to the same product?

Which is a common trap when interpreting enthalpy change values?

What is the significance of the term 'standard state' in thermochemistry?

In the combustion of benzene, how much energy is liberated per mole?

Which of the following reactions generally indicates a decrease in enthalpy?

What is the bond enthalpy of a C-H bond in methane if the total enthalpy change for the breaking of all bonds is +1665 kJ/mol?

How is the standard enthalpy of a reaction related to bond enthalpies?

What is the correct expression for calculating the enthalpy change of a reaction using Kirchhoff's equation?

What is the Gibbs free energy equation?

If a reaction has a negative ∆G, what can be concluded about its spontaneity?

Under what condition is a reaction with positive ∆H and positive ∆S spontaneous?

Which of the following statements about spontaneous processes is true?

Which term describes the measure of disorder in a system?

What does a positive value of ∆G imply for a reaction?

When considering spontaneity, which factor directly influences Gibbs free energy the most?

Which pair of thermodynamic quantities must be evaluated to determine spontaneity?

Which scenario would likely favor the spontaneity of a reaction?

What role does temperature play in reactions that are both endothermic and result in increased entropy?

Which thermodynamic principle asserts that entropy of an isolated system tends to increase over time?

What happens to the spontaneity of a reaction as temperature increases if both enthalpy and entropy are positive?

Which of the following best describes a reaction at equilibrium considering Gibbs energy?

What does the term 'spontaneous' signify in a chemical reaction context?

What does a negative Gibbs free energy change (∆G) indicate about a reaction?

Which of the following equations relates Gibbs free energy change to the equilibrium constant?

At equilibrium, what is the value of ∆G for a reversible reaction?

How does an increase in temperature typically affect the Gibbs free energy for an endothermic reaction?

In a reaction where both ∆H and ∆S are negative, under what condition is the reaction spontaneous?

Which factor does NOT affect the Gibbs free energy change of a reaction?

If ∆G > 0 for a reaction under standard conditions, what can be concluded?

A reaction's equilibrium constant decreases as the temperature increases. Which statement is true about the reaction?

For a reaction characterized by ∆H > 0 and ∆S > 0, when will the reaction be spontaneous?

If the standard Gibbs free energy change of a reaction is given as ∆Gₐ = -RT ln K, what can be inferred about the formation of products?

Consider a reversible reaction at equilibrium. If the concentration of products is increased, how does this affect ∆G?

Which of the following statements is valid regarding ∆G and reaction spontaneity?

In thermodynamics, what does the term 'spontaneity' refer to?

If a reaction has ∆H < 0 and ∆S < 0, under what temperature ranges will the reaction be spontaneous?

What is the significance of a reaction where ∆G = 0?

Single Answer MCQ

Q-00054863

View explanation

Thermodynamics Practice Worksheets

Download and practice Thermodynamics worksheets to improve problem-solving accuracy and speed for CBSE Class 11 Chemistry exams.

Thermodynamics - Practice Worksheet

This worksheet covers essential long-answer questions to help you build confidence in Thermodynamics from Chemistry Part - I for Class 11 (Chemistry).

Practice

Questions

Explain the concept of a system and its surroundings in thermodynamics. Provide examples of open, closed, and isolated systems.

In thermodynamics, a system is defined as a specific part of the universe that we focus on for analysis, whereas everything else is termed as the surroundings. An open system allows both energy and matter to exchange with surroundings, like a pot of boiling water. A closed system can exchange energy but not matter, like a sealed container of gas. An isolated system cannot exchange either energy or matter, such as thermos flasks. The universe is the sum of the system and surroundings.

What is internal energy, and how does it change in a system? Discuss the contributions of heat and work to internal energy changes.

Internal energy (U) is the total energy contained in a system, including kinetic and potential energy of the molecules. It changes when heat is added or removed (q) and when work is done on or by the system (w). According to the first law of thermodynamics, the change in internal energy is given by ΔU = q + w. This equation emphasizes that internal energy is a state function, depending only on initial and final states, not on the path taken.

Define and differentiate between extensive and intensive properties. Provide two examples of each type.

Extensive properties depend on the amount of substance present, such as mass and volume. Intensive properties do not depend on the quantity, like temperature and pressure. For example, mass and volume are extensive because they change when you have more or less of the material, whereas temperature and pressure remain the same regardless of the amount of substance. This distinction is crucial in thermodynamic processes.

Express the first law of thermodynamics mathematically and explain its significance in thermodynamics.

The first law of thermodynamics states that the total energy of an isolated system is constant. Mathematically, it is expressed 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. This law signifies the principle of conservation of energy; energy can neither be created nor destroyed, only transformed from one form to another.

What is enthalpy (H)? Describe its relationship with internal energy and pressure-volume work.

Enthalpy (H) is a thermodynamic quantity defined as H = U + pV, where U is the internal energy, p is pressure, and V is volume of the system. The change in enthalpy (ΔH) during a process is related to heat transfer at constant pressure. Enthalpy takes into account not only internal energy but also the work done by or on the system during expansion or compression.

Explain Hess's Law of constant heat summation and provide an example of its application.

Hess's Law states that the total enthalpy change in a chemical reaction is the same, no matter how many steps the reaction takes. This law is based on the fact that enthalpy is a state function. For example, if we know the enthalpy changes for a series of reactions leading to the same products, we can sum them to find the overall enthalpy change. An example is finding the enthalpy change for the formation of CO from its elements through intermediate reactions like C + O2 → CO2 and CO2 + C → 2CO.

What is entropy (S) and how does it relate to spontaneity in thermodynamic processes?

Entropy (S) is a measure of the disorder or randomness in a system. It quantifies how much energy in a system is not available for doing work. The second law of thermodynamics states that for spontaneous processes, the total entropy change (system + surroundings) must be greater than zero (ΔS_total > 0). This means that spontaneous processes lead to an increase in disorder in the universe.

Define Gibbs free energy (G) and explain its significance in predicting the spontaneity of reactions.

Gibbs free energy (G) is a thermodynamic potential that measures the maximum reversible work obtainable from a thermodynamic system at constant temperature and pressure. It is defined as G = H - TS, where H is enthalpy, T is temperature, and S is entropy. A reaction is spontaneous if the change in Gibbs free energy (ΔG) is negative (ΔG < 0). This means that the reaction can occur without the input of work.

How do temperature and enthalpy changes affect the spontaneity of a reaction? Discuss with examples.

The spontaneity of a reaction depends both on enthalpy change (ΔH) and entropy change (ΔS), as incorporated in the Gibbs free energy equation (ΔG = ΔH - TΔS). If ΔG is negative, the reaction is spontaneous. For example, exothermic reactions (negative ΔH) tend to be spontaneous at all temperatures. In contrast, endothermic reactions (positive ΔH) can become spontaneous at high temperatures if the entropy change is positive and large enough to outweigh the enthalpy term.

Thermodynamics - Mastery Worksheet

This worksheet challenges you with deeper, multi-concept long-answer questions from Thermodynamics to prepare for higher-weightage questions in Class 11.

Mastery

Questions

Explain the first law of thermodynamics and illustrate how it can be applied to calculate the internal energy change for a process involving both work and heat transfer. Include examples of different types of systems.

The first law of thermodynamics states that the energy of an isolated system is constant. It can be expressed mathematically 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. For example, in a closed container where heat is added (q > 0) and work is done on the system (w > 0), both contribute to an increase in internal energy (ΔU > 0).

Compare and contrast the concepts of internal energy (U) and enthalpy (H). Under what conditions are they equivalent, and when would one be preferred over the other in thermodynamic calculations?

Internal energy (U) is the total energy of a system, encompassing kinetic and potential energies of molecules, while enthalpy (H = U + pV) incorporates the pressure-volume work that can be done. They are equivalent when processes occur at constant volume and are considered in situations where pressure changes, such as at constant atmospheric pressure. For reactions happening at constant pressure, H is preferred for calculations involving heat transfer.

Describe Hess’s law and provide a practical example demonstrating how it can be used to calculate enthalpy changes for a reaction that cannot be measured directly.

Hess’s law states that the total enthalpy change for a reaction is the sum of the enthalpy changes for individual steps leading to the same final state. For example, to calculate the enthalpy change for the combustion of carbon, we can use the enthalpy changes of the combustion of related compounds. If the combustion of CO2 gives -393.5 kJ/mol, factored with the combustion of C and O2 can yield the overall reaction enthalpy.

What is entropy? Discuss its significance in thermodynamics and explain how it relates to the second law of thermodynamics with examples.

Entropy (S) is a measure of disorder or randomness in a system. It quantifies the number of microscopic configurations that correspond to a macroscopic state. The second law of thermodynamics states that in an isolated system, the total entropy always increases over time, indicating that natural processes tend to move toward a state of maximum disorder. For instance, the melting of ice (solid to liquid) is accompanied by an increase in entropy.

Define Gibbs free energy and explain its role in predicting spontaneity of reactions. How does it relate to enthalpy and entropy?

Gibbs free energy (G) is defined by the equation G = H - TS. It combines the concepts of enthalpy and entropy to predict the spontaneity of reactions at constant temperature and pressure. If ΔG is negative, the process is spontaneous; if positive, it is non-spontaneous. For example, a reaction with ΔH < 0 and ΔS > 0 will always be spontaneous.

Calculate the change in enthalpy for the endothermic reaction absorbing heat at constant pressure and relate it to internal energy changes. Include a calculation for heat capacity in your response.

The enthalpy change (ΔH) in an endothermic reaction that absorbs q amount of heat at constant pressure can be directly related to the internal energy change as ΔU = ΔH - pΔV. For a process involving specific heat capacity (C), where C = q/ΔT, we can calculate ΔH = C * ΔT. For instance, if 100 J of heat is absorbed at a constant pressure with C = 4.18 J/g°C and ΔT is 25°C, ΔH = C * ΔT = 4.18 * 25 = 104.5 J.

Investigate the differences between spontaneous and non-spontaneous processes using examples that show the role of energy changes in determining reaction feasibility.

Spontaneous processes, such as the rusting of iron or combustion of fuels, occur naturally without continuous external input, typically indicating a negative Gibbs free energy change. Non-spontaneous processes, like the conversion of diamond to graphite, require energy input. In spontaneous reactions, enthalpy often decreases while entropy increases; they can be quantified with ΔG calculations to establish their feasibility.

Using the relationship between ΔG and K, calculate the equilibrium constant from given Gibbs energy changes.

The relationship given by ΔG° = -RT ln(K) allows the calculation of equilibrium constant K from Gibbs free energy change ΔG°, where R is the universal gas constant (8.314 J/K∙mol) and T is in Kelvin. For example, if ΔG° is -40 kJ/mol, converting to J gives K = e^(-ΔG°/RT). Substituting these values helps find K.

Thermodynamics - Challenge Worksheet

The final worksheet presents challenging long-answer questions that test your depth of understanding and exam-readiness for Thermodynamics in Class 11.

Challenge

Questions

Evaluate the implications of the First Law of Thermodynamics in the context of a closed system undergoing a chemical reaction, where heat is neither absorbed nor released.

Discuss how energy conservation is maintained in the reaction and provide examples of possible work done and changes in internal energy.

Critically analyze how the concept of enthalpy changes during a phase transition, such as melting ice, affects the overall energy balance of the system.

Explain the relationship between enthalpy, temperature, and the irreversibility of the process using specific heat capacities.

Discuss the significance of Hess's law in predicting the enthalpy change for a reaction that cannot be easily measured directly.

Provide examples of reactions where Hess's law can be applied and demonstrate calculations involving multiple steps.

Evaluate the relationship between Gibbs free energy and spontaneity of reactions at constant temperature and pressure.

Analyze various scenarios with different signs of Gibbs free energy and relate them to spontaneous and non-spontaneous processes.

Analyze a real-life scenario in which a chemical process with a positive enthalpy change can still be spontaneous under specific conditions.

Apply the Gibbs energy equation to explain such conditions and provide contextual examples.

Examine the concept of standard state and how it applies to the calculation of enthalpy changes for chemical reactions.

Illustrate your points with examples of substances in their standard states and calculations of their enthalpy changes.

Evaluate the role of entropy as a criterion for spontaneity in thermodynamic processes, despite the First Law of Thermodynamics.

Discuss how changes in entropy can drive processes where energy conservation does not alone dictate the outcome.

Assess the significance of pressure-volume work in processes involving gases and its contribution to the internal energy change of a system.

Provide derivations for work done in different types of expansion (isothermal, adiabatic) and relate them to energy conservation.

Describe how temperature influences the spontaneity of reactions that are both exothermic and endothermic.

Create a comprehensive overview outlining how Gibbs free energy varies with temperature changes.

Investigate the relationship between reaction coordinates and the potential energy surface concerning Gibbs energy and spontaneity.

Delve into the interpretation of energy profiles for reactions, including transition states and their Gibbs energy implications.

Thermodynamics Formula Sheet

Use this Class 11 Chemistry Thermodynamics Formula Sheet for quick revision before school exams and CBSE exams. It brings together the important formulas, key concepts, and worked examples in one place so students can revise faster and download a printable PDF for offline study.

Important Formulas

∆U = q + w

∆U is the change in internal energy, q is the heat added to the system, and w is the work done on the system. This equation expresses the first law of thermodynamics, stating energy conservation.

q = C × ∆T

q is the heat transferred, C is the heat capacity, and ∆T is the change in temperature. This formula calculates the heat absorbed or released when the temperature of a substance changes.

∆H = ∆U + ∆n_g RT

∆H is the change in enthalpy, ∆U is the change in internal energy, ∆n_g is the change in moles of gas, R is the universal gas constant, and T is the temperature in Kelvin. This relates enthalpy change to changes in states involving gases.

q_p = ∆H

q_p is the heat at constant pressure, which equals the change in enthalpy (∆H). This is useful for reactions occurring at a constant atmospheric pressure.

w = - p_ex ∆V

w is the work done on or by the system, p_ex is the external pressure, and ∆V is the change in volume. This equation calculates work in expansion (or compression) processes.

∆S = q_rev / T

∆S is the change in entropy, q_rev is the reversible heat transfer, and T is the temperature in Kelvin. This formula defines how the entropy of a system changes with heat transfer at a constant temperature.

∆S_total = ∆S_system + ∆S_surroundings

This equation expresses the second law of thermodynamics, stating that the total change in entropy (∆S_total) is the sum of changes in entropy of the system and its surroundings.

G = H - TS

G is the Gibbs free energy, H is the enthalpy, T is the temperature in Kelvin, and S is the entropy. This equation is central to determining spontaneity in chemical processes.

∆G = ∆H - T∆S

∆G is the change in Gibbs free energy, ∆H is the enthalpy change, and ∆S is the entropy change. This relationship allows us to evaluate the spontaneity of a reaction.

∆G = -RT ln K

Where K is the equilibrium constant. This equation relates the standard Gibbs free energy change to the equilibrium constant of a reaction at a given temperature.

Worked Examples

PV = nRT

This is the ideal gas law, where P is pressure, V is volume, n is the number of moles of gas, R is the ideal gas constant, and T is temperature in Kelvin. It relates the state of an ideal gas.

H = U + PV

This defines enthalpy (H) as the sum of internal energy (U) and the product of pressure (P) and volume (V). It is essential for understanding enthalpy in reactions.

∆H = ΣH_products - ΣH_reactants

This is Hess's law for adding enthalpy changes in reactions. It states that the enthalpy change for a reaction is the sum of the enthalpy changes for each step.

q = n∆H_f

Where q is the heat absorbed/released, n is the number of moles, and ∆H_f is the heat of formation. This applies in calculating heat during formation reactions.

p∆V = ∆n_g RT

This equation relates the pressure and volume change (∆V) to the change in number of moles of gas (∆n_g), useful in gas expansion/compression calculations.

K = e^(-∆G/RT)

The relationship between the equilibrium constant (K), Gibbs free energy change (∆G), the gas constant (R), and temperature (T). It indicates how free energy relates to equilibrium position.

∆U = q + w

This is the first law of thermodynamics restated, where ∆U is the change in internal energy. Typically used for closed systems to understand energy transfers.

q = m×c×∆T

Where m = mass, c = specific heat capacity, and ∆T = temperature change. This formula calculates heat transfer in a specific heat scenario.

∆H = ∆U + ∆nRT

Expresses the enthalpy change of a system as a function of the internal energy change plus the product of number of moles of gas change, the temperature, and the gas constant.

∆S = k*ln(Ω)

Where k is Boltzmann's constant and Ω represents the number of microscopic configurations that correspond to a macroscopic state. It connects statistical mechanics with thermodynamics.

Explore More Thermodynamics Resources

Explore more chapter resources to strengthen your understanding and prepare for exams.

Practice Questions Worksheets Revision Notes Flashcards Download PDFs

Thermodynamics Frequently Asked Questions

Explore the principles of thermodynamics, encompassing energy transformations in chemical processes, internal energy, enthalpy changes, and the significance of Gibbs energy in determining reaction spontaneity.

QWhat is a thermodynamic system?

A thermodynamic system is a specific region of the universe selected for analysis, bounded by a boundary that can be real or imaginary. It includes everything contained within that boundary.

QWhat are the three types of thermodynamic systems?

Thermodynamic systems can be classified as open systems, where both matter and energy can exchange with the surroundings; closed systems, where only energy can be exchanged; and isolated systems, where neither matter nor energy can exchange with the surroundings.

QWhat is internal energy?

Internal energy (U) is the total energy contained within a system, encompassing kinetic and potential energy of molecules. It changes with heat transfer, work done on the system, or mass transfer.

QHow is work defined in thermodynamics?

In thermodynamics, work is defined as the energy transfer associated with a process that involves a system changing volume against an external pressure, often quantified using the formula w = -p_ex ΔV.

QWhat is the first law of thermodynamics?

The first law of thermodynamics states that energy cannot be created or destroyed in an isolated system, meaning the total energy change of a system equals the heat added to the system minus the work done by the system.

QWhat is enthalpy?

Enthalpy (H) is a thermodynamic state function defined as the internal energy of a system plus the product of its pressure and volume (H = U + pV). It reflects the heat content of a system at constant pressure.

QHow can enthalpy changes be calculated?

Enthalpy changes can be calculated using the formula ΔH = ΔU + Δ(nRT), where ΔU is the change in internal energy, n is the moles of gases involved, and R is the universal gas constant.

QWhat is the significance of Gibbs free energy?

Gibbs free energy (G) is a crucial concept that combines enthalpy and entropy to predict spontaneity in chemical processes. A reaction is spontaneous when ΔG is negative, indicating a net release of useful energy.

QWhat does it mean for a process to be spontaneous?

A spontaneous process occurs without the need for continuous external intervention, often associated with a decrease in free energy or an increase in entropy in the system, leading to a natural tendency to occur.

QWhat is Hess's Law?

Hess's Law states that the total enthalpy change for a reaction is the same, regardless of whether the reaction occurs in one step or multiple steps. This property allows the calculation of enthalpy changes using other known values.

QWhat is the relationship between temperature and spontaneity?

Temperature affects spontaneity by influencing the entropy change in processes involving heat transfer. Reactions with positive entropy changes may become spontaneous at higher temperatures, while others may not.

QWhat happens to entropy in a spontaneous reaction?

In a spontaneous reaction, the total entropy of the system and surroundings typically increases. This rise in entropy indicates a greater degree of disorder and supports the favorability of the reaction.

QHow do phase changes relate to thermodynamics?

Phase changes involve energy transformations that can be analyzed using thermodynamic principles. Each phase transition, such as melting or vaporization, has associated enthalpy changes that reflect the energy required for the change.

QWhat factors influence the enthalpy change during a reaction?

The enthalpy change during a reaction is influenced by the nature of the reactants, the products formed, temperature, pressure, and the state of substances involved (solid, liquid, or gas).

QCan you provide an example of a system and its surroundings?

In a chemical reaction conducted in a beaker, the beaker and its contents represent the system. The room's air, the table, and the surrounding environment constitute the surroundings, which can interact with the system.

QWhat is the difference between extensive and intensive properties?

Extensive properties depend on the amount of matter present in a system, such as mass or volume, while intensive properties do not, examples being temperature and pressure.

QHow is the heat capacity of a substance related to temperature change?

Heat capacity denotes the amount of heat required to change a substance's temperature. It is crucial for calculating temperature changes via q = C∆T, where q is the heat exchanged and C is heat capacity.

QWhat does a negative enthalpy change indicate?

A negative enthalpy change indicates that a reaction is exothermic, meaning it releases heat to the surroundings, typically causing the temperature of the surroundings to rise.

QWhat is the significance of standard enthalpy of formation?

The standard enthalpy of formation is the change in enthalpy when one mole of a compound is formed from its elements under standard conditions. It provides a reference point for calculating energy changes in various reactions.

QWhy does the dissolution of some ionic compounds require energy?

The dissolution of ionic compounds may require energy due to strong ionic bonds needing to be broken. If the lattice enthalpy is high, it may be endothermic, leading to positive enthalpy changes during dissolution.

QHow is calorimetry used in thermodynamics?

Calorimetry is a method used to measure the heat changes associated with chemical or physical processes. It involves using calorimeters to quantify heat transfer by monitoring temperature changes, under controlled conditions.

QWhat is the relation between bond enthalpy and thermodynamics?

Bond enthalpy is the energy required to break chemical bonds. In thermodynamics, it helps calculate the enthalpy change of reactions by comparing the bond enthalpies of reactants and products, providing insights into energy changes.

QWhat does it mean if a reaction has a positive Gibbs free energy?

A positive Gibbs free energy indicates a reaction is not spontaneous under the given conditions, meaning it requires external energy input to proceed. It suggests that the enthalpy change is greater than the temperature times the entropy change.

QWhy do reactions reach equilibrium?

Reactions reach equilibrium when the forward and reverse reaction rates are equal, leading to constant concentrations of reactants and products. At this point, the Gibbs free energy change is zero, indicating a stable state.

Reviewed & Verified by Edzy Expert

Academically Reviewed

Dr. Meera Nair

CBSE Physics & Chemistry Reviewer

Senior Physics and Chemistry educator with 15+ years of experience in CBSE senior secondary curriculum and competitive exam content.

Ph.D. ChemistryM.Sc. Physics

Thermodynamics PDF Downloads

Download worksheets, revision guides, formula sheets, and the official textbook PDF for Thermodynamics.

Thermodynamics Official Textbook PDF

Download the official NCERT/CBSE textbook PDF for Class 11 Chemistry.

Official PDFEnglish EditionNCERT Source

Download PDF

Thermodynamics Revision Guide

Use this one-page guide to revise the most important ideas from Thermodynamics.

Best for1-page chapter recap

Download PDF

Thermodynamics Formula Sheet

Download the Thermodynamics formula sheet PDF with important formulas, worked examples, and quick revision support for exam preparation.

Best forImportant formulas for quick revision

Download Formula Sheet PDF

Thermodynamics Practice Worksheet

Solve basic and application-based questions from Thermodynamics.

Best forCore practice set

Download PDF

Thermodynamics Mastery Worksheet

Work through mixed Thermodynamics questions to improve accuracy and speed.

Best forMixed difficulty set

Download PDF

Thermodynamics Challenge Worksheet

Try harder Thermodynamics questions that test deeper understanding.

Best forFor deeper problem solving

Download PDF

Thermodynamics Question Bank

Download important questions and exam-style prompts from Thermodynamics.

Best forPrintable question set

Download PDF

Thermodynamics Flashcards

Revise key terms and definitions from Thermodynamics with interactive flashcards. Quick recall practice for CBSE Class 11 Chemistry.

These flash cards cover important concepts from Thermodynamics in Chemistry Part - I for Class 11 (Chemistry).

1/19

What is thermodynamics?

1/19

Thermodynamics is the study of energy transformations and the laws governing energy changes in physical and chemical processes.

How well did you know this?

Not at allPerfectly

2/19

What is a system in thermodynamics?

2/19

A system is the part of the universe that is being studied, while the surroundings are everything outside the system.

How well did you know this?

Not at allPerfectly

Active

3/19

What are closed, open, and isolated systems?

Active

3/19

Open systems can exchange both energy and matter; closed systems exchange only energy; isolated systems exchange neither.

How well did you know this?

Not at allPerfectly

4/19

What is internal energy?

4/19

Internal energy (U) is the total energy contained within a system, resulting from kinetic and potential energies of particles.

5/19

State the first law of thermodynamics.

5/19

The first law states that energy cannot be created or destroyed, only transformed: ∆U = Q - W (change in internal energy equals heat added minus work done).

6/19

How is work defined in thermodynamics?

6/19

Work is the energy transfer that occurs when a force is applied over a distance, often due to volume change against pressure in a system.

7/19

What is heat (Q) in thermodynamics?

7/19

Heat is the energy transferred between a system and its surroundings due to a temperature difference.

8/19

What is enthalpy?

8/19

Enthalpy (H) is a thermodynamic property defined as H = U + PV, representing the total heat content of a system at constant pressure.

9/19

What are state functions?

9/19

State functions are properties that depend only on the state of a system, not how it got there, such as internal energy (U) and enthalpy (H).

10/19

What is the difference between ∆U and ∆H?

10/19

∆U represents the change in internal energy, while ∆H represents the change in enthalpy, which includes the effect of pressure-volume work.

11/19

How is enthalpy change calculated for reactions?

11/19

Enthalpy change (∆H) can be calculated using calorimetry or from standard enthalpies of formation.

12/19

What is Hess’s law?

12/19

Hess's law states that the total enthalpy change for a reaction is the sum of the enthalpy changes for individual steps, regardless of the pathway taken.

13/19

What is the difference between extensive and intensive properties?

13/19

Extensive properties depend on the amount of substance (e.g., mass, volume), while intensive properties do not (e.g., temperature, pressure).

14/19

Define spontaneous processes.

14/19

Spontaneous processes are natural processes that occur without external intervention, often associated with a decrease in free energy.

15/19

What is entropy?

15/19

Entropy is a measure of disorder or randomness in a system, with higher entropy indicating greater disorder.

16/19

What is Gibbs free energy?

16/19

Gibbs free energy (G) is a thermodynamic potential that indicates the spontaneity of a process: ∆G = ∆H - T∆S.

17/19

How is ∆G related to spontaneity?

17/19

A negative ∆G indicates a spontaneous process, while a positive ∆G indicates a non-spontaneous process.

18/19

What is the relationship between ∆G and equilibrium constant (K)?

18/19

The relationship is given by the equation ∆G = -RT ln(K), where R is the gas constant and T is temperature in Kelvin.

19/19

What are common misconceptions in thermodynamics?

19/19

Common mistakes include confusing heat with temperature, misunderstanding work as energy, and not recognizing that heat is energy transfer.

Practice Thermodynamics with Interactive Duels

Live Academic Duel

Master Thermodynamics via Live Academic Duels

Challenge your classmates or test your individual retention on the core concepts of CBSE Class 11 Chemistry (Chemistry Part - I). Compete in speed-recall question rounds matched explicitly to the latest syllabus milestones for Thermodynamics.

CBSE-aligned questions

Instant speed-recall rounds

Quick, competitive practice on Thermodynamics with zero setup.