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Thermodynamics - One Shot | Class 11 Chemistry | JEE 2024 | Nabamita Ma'am

Vedantu JEE English・160 minutes read

The session focuses on thermodynamics, discussing three laws, system definitions, types of systems, thermodynamic processes, internal energy, and the first law of thermodynamics. It also covers isothermal, isobaric, isochoric processes, work, and heat transfer, as well as entropy, Gibbs free energy, and the second and third laws of thermodynamics, aiming to explain the feasibility and extent of chemical reactions.

Insights

Understanding thermodynamics is crucial for predicting the feasibility and extent of chemical reactions, making it essential for JW and NEET exams.
The laws of thermodynamics focus on energy changes in macroscopic systems, excluding microscopic environments.
Thermodynamics has limitations, such as an inability to predict reaction paths or speeds and its inapplicability to microscopic systems.
Definitions of system, surrounding, and universe are explained, with the universe comprising the sum of the system and surrounding.
Internal energy, a crucial concept in thermodynamics, changes due to matter exchange, heat transfer, or work done on or by the system, affecting the system's total energy.

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Recent questions

What is the focus of thermodynamics?
Understanding heat and system changes.
What are the disadvantages of thermodynamics?
Not predicting reaction path or speed.
What are the three laws of thermodynamics?
Laws focusing on energy changes in systems.
What is the definition of an open system?
System allowing exchange of matter and energy.
What is the significance of state functions?
Values dependent on system's state, not path.

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Summary

00:00

"Thermodynamics Essentials for JW and NEET"

The session is about studying the chapter of thermodynamics with the teacher Nomita Bach on the Vidant English Channel.
The focus is on understanding thermodynamics, which deals with heat and changes in a system.
The chapter is crucial for JW and NEET exams, as it helps predict the feasibility and extent of chemical reactions.
Three laws of thermodynamics are discussed, focusing on macroscopic systems.
The laws deal with energy changes in macroscopic systems, not microscopic ones.
Thermodynamics has disadvantages, such as not predicting the path or speed of a reaction and not being applicable to microscopic environments.
Definitions of system, surrounding, and universe are explained, with the universe being the sum of the system and surrounding.
Different types of systems are introduced, with an open system allowing the exchange of matter and energy between the system and surroundings.
The concept of adiabatic walls, also known as diathermal walls, is mentioned as allowing the exchange of matter and energy.
The session is expected to last 3 to 4 hours, starting at 7:00 PM, covering the entire chapter of thermodynamics.

16:28

Thermodynamics: Systems, Processes, and Equations

Closed system: Putting a lid on a coaster shows no exchange of matter but energy exchange occurs.
Diagram: Illustrating a closed system with a lid and mentioning the absence of matter exchange but energy exchange.
Isolated system: Describing an isolated system with no exchange of energy or matter between system and surroundings.
Thermos flask: Explaining how a thermos flask maintains water temperature for a day or two, introducing the concept of an isolated system.
Types of thermodynamic processes: Detailing isothermal, isobaric, isochoric, and adiabatic processes.
Cyclic process: Defining a cyclic process where initial and final states are the same, emphasizing the formation of a loop in the graph.
Wall types: Differentiating between adiabatic and diathermal walls based on the system's requirements.
State functions: Defining state functions as values dependent on the system's state, not on how it was reached.
Path functions: Comparing path functions to weight loss journeys, where the process matters and is discussed.
Gibbs-Helmholtz equation: Using the equation to explain state and path functions, highlighting the difference in approach and interest between the two.

34:25

"Thermodynamic Properties and Energy Changes Explained"

Achieving a rank one in JEE is of high interest to many, unlike achieving a 75% criteria in class 12.
State functions, like achieving a rank one, are of significant interest, while 75% criteria is not as noteworthy.
Examples of state functions include pressure, volume, temperature, Delta G, Delta H, and Delta S.
Path functions are those that depend on the path taken to reach a specific value, such as work and heat transfer.
Extensive properties, like weight and volume, depend on the quantity or size of matter present in the system.
Intensive properties, like temperature and refractive index, do not depend on the quantity or size of matter.
Internal energy is the sum of all energies within a thermodynamic system and is an extensive property.
The absolute value of internal energy cannot be determined, but the change in internal energy can be calculated.
Internal energy is denoted by U or E and changes when matter enters or leaves the system, heat passes into or out of the system, or work is done on or by the system.
The total energy of the system remains constant, with Delta U being negative if energy is released, positive if absorbed, and zero for an isothermal process.

55:53

Understanding Internal Energy Changes in Thermodynamics

Isothermal process: Delta U is zero
Energy release by system leads to negative Delta U
Energy absorption by system results in positive Delta U
Internal energy change by work in adiabatic system involves mechanical or electrical work
Change in state of system remains the same regardless of path taken
Internal energy change by heat involves heat transfer, with negative Q when system releases heat and positive Q when system absorbs heat
Internal energy change by heat and work is represented by Delta U = Q + W, with Delta U as a state function and Q + W as a path function
First law of thermodynamics states energy of an isolated system is constant
Internal energy change by pressure volume work involves gas expansion, with pressure or volume work done leading to changes in internal energy
Introduction to reversible and irreversible processes through understanding of pressure volume work and gas expansion.

01:16:03

"Piston Movement and Work Calculation in Thermodynamics"

External pressure is exerted on the piston, causing it to move inside.
Volume change is determined by the length multiplied by the area of the cross-section.
Pressure (P) is calculated as force divided by area.
The force on the piston is equal to the external pressure multiplied by the area.
Work done is calculated as P external multiplied by the change in volume (Delta V).
Compression results in negative Delta V, indicating a decrease in volume.
For work done to be positive, external pressure must be negative.
Single-step changes differ from gradual changes, requiring integration for smaller increments.
Reversible processes involve gradual changes, while irreversible processes involve single-step changes.
Work done in a reversible process is calculated as -2.303 nRT log(V2/V1).

01:32:55

"Thermodynamics: Reversible vs Irreversible Processes"

Studying for half an hour daily can be tedious due to waiting for class and counting pages.
After completing thermodynamics, starting a new chapter the next day is possible.
Reversible expansion requires more work than irreversible expansion.
Reversible processes are imaginary, while irreversible processes are natural.
Reversible processes maintain equilibrium throughout, while equilibrium exists only at the beginning and end in irreversible processes.
Maximum work is obtained in reversible processes, while work done is minimum in irreversible processes.
Reversible processes follow a reversible path and take infinite time, while irreversible processes follow a finite path.
In free expansion of gas with zero external pressure, work done is also zero.
For an isothermal change, Delta U is zero, as there is no heat exchange.
Graphs for isothermal, isobaric, and isochoric processes can be determined using the ideal gas law PV = nRT.

01:58:07

Temperature Conversion, Mass Calculation, Gas Expansion, Thermodynamic Changes

The text discusses a calculation involving the conversion of temperature from Kelvin to another unit.
It mentions the calculation of mass based on the given temperature and a specific formula.
The result of the mass calculation is determined to be 58.5 kg.
Work done in the free expansion of gas is briefly touched upon.
The text delves into the concept of isothermal irreversible change, detailing the equation Q = -W = P external (V final - V initial).
It contrasts isothermal irreversible change with isothermal reversible change, explaining the equation Q = -W = 2.303nRT log (V final / V initial).
The concept of adiabatic change is introduced, with the equation Q = 0 and Delta U = w ad.
The sign convention of work and heat in the first law of thermodynamics is explained, highlighting the positivity of work done on the system and the negativity of work done by the system.
The text emphasizes the distinction between exothermic and endothermic reactions in relation to Delta H being negative or positive.
Enthalpy is defined as internal energy plus pressure multiplied by volume, with the significance of the difference between Delta H and Delta U discussed in the context of solids, liquids, and gases.
The significance of the difference between Delta H and Delta U is noted to be more pronounced in systems involving gases, leading to a discussion on the enthalpy of gases.

02:18:26

Gas Enthalpy and Heat Capacity Relations

Enthalpy of gases is discussed, focusing on the ideal gas equation PV = nRT for reactants and products at constant pressure.
The equation is manipulated to show that Delta H = Delta U + Delta nRT, emphasizing constant pressure conditions.
Heat capacity is defined as the amount of heat needed to raise the temperature of a given mass by one degree Celsius, denoted as C subscript T.
Specific heat capacity is explained as the heat required to raise the temperature of unit mass by one degree Celsius, denoted as C.
The relationship between constant pressure (CP) and constant volume (CV) in an ideal gas is explored, leading to the conclusion that CP - CV = R.
Formulas for molar heat capacity at constant pressure (CPM), specific heat capacity at constant pressure (CPS), and heat capacity at constant volume (CV) are detailed.
Important relations for isothermal, isobaric, isochoric, and adiabatic processes are outlined, including Mayor's relation (CPM - CVM = R).
Heat capacity is identified as a path function, while CP and CV are constants and not path functions.
The ratio CP/CV is highlighted as equal to R, representing the degree of freedom in the system.
The ratio gamma (γ) is defined as CP/CV, with a mention of a table for further reference.

02:38:02

Gas Degrees of Freedom and Enthalpy Calculation

Gamma formula: 1 + 2/F
Degree of Freedom (F) explained as the degree of movement.
Monoatomic, diatomic, and triatomic gases discussed.
Monoatomic gas has 3 degrees of freedom.
Diatomic gas has 5 degrees of freedom.
Triatomic gas has 6 degrees of freedom.
Calculation of specific heat capacities for different gases.
Introduction to reaction enthalpy and its significance.
Standard enthalpy of reaction defined as Delta H with a circle.
Explanation of standard enthalpy of fusion, vaporization, and sublimation.

02:56:27

Enthalpy in Chemistry: Key Concepts Explained

Standard enthalpy of fusion is the amount of heat required to convert one mole of a solid substance to a liquid in its standard state, an endothermic reaction.
Standard enthalpy of vaporization is the heat needed to vaporize one mole of a liquid in its standard state, also an endothermic reaction.
Standard enthalpy of sublimation is the change in enthalpy when one mole of a solid substance sublimes at a standard state, another endothermic reaction.
Enthalpy change in reactions is determined by intermolecular interactions.
Standard enthalpy of formation is the enthalpy change for forming one mole of a compound from its elements in their reference state.
Standard enthalpy of formation is a special case of reaction enthalpy where compounds form only from constituent elements.
Thermochemical equations are balanced chemical reactions with enthalpy values indicating the number of moles of substances involved.
Coefficients in thermochemical equations represent the number of moles, not molecules.
Hess's Law states that the standard reaction enthalpy of a multi-step reaction is the sum of all intermediate enthalpies.
Different types of standard enthalpies include combustion, atomization, bond enthalpy, and lattice enthalpy. Bond Haber cycle is a method to calculate lattice enthalpy.

03:16:40

"Chemical Reactions and Thermodynamics Explained"

Dilution involves adding more solvent to a solution to create a dilute solution.
Enthalpy of solution is the enthalpy change when one mole of a substance dissolves in a specified amount of solvent.
NaCl in a solution dissociates into Na+ and Cl- ions.
Enthalpy of dilution is the difference in enthalpy values between different dilutions of a substance.
Enthalpy of dilution is the heat withdrawn from the surroundings when additional solvent is added to a solution.
Spontaneity in reactions is determined by entropy, which measures the degree of randomness or disorder in a system.
Solids have minimum entropy, while gases have maximum entropy.
Change in entropy is directly proportional to heat added to the system and inversely proportional to temperature.
For reversible processes, the change in entropy of the universe is zero, while for irreversible processes, it is greater than zero.
Gibbs free energy (G) is used to determine the spontaneity of a reaction, calculated as G = H - TS or ΔG = ΔH - TΔS using the Gibbs-Helmholtz equation.

03:38:50

Determining Spontaneity in Reactions Using Equations

Gibbs Helmholtz equation is used to determine spontaneity in reactions.
A negative Delta G indicates a spontaneous reaction, while a positive Delta G indicates a non-spontaneous reaction.
The second law of thermodynamics states that in an exothermic reaction, heat release increases disorder in the surroundings, leading to a positive entropy change and a spontaneous reaction.
The third law of thermodynamics asserts that the entropy of a pure crystalline substance approaches zero as the temperature nears absolute zero.
The mathematical expression for the third law of thermodynamics involves integrating entropy from 0 to T CP DT divided by T.
A chart detailing different types of processes and their graphs is provided for further understanding.

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