Sections in the chapter (Jauhar’s Book CBSE)
4.1 Some basic terms and concepts
4.2 Internal energy, heat and work
4.3 First law of thermodynamics
Practice Problems 4.1 to 4.6
4.4 Enthalpy and Enthalpy change
P.P 4.7 to 4.10
4.5 Limitations of first law of thermodynamics: Need of second law
4.6 Randomness and spontaneity
4.7 Entropy
4.8 Second law of thermodynamics
4.9 Entropy change during phase transition
P.P. 4.11 to 4.17
4.10 Gibbs Free energy and free energy change
4.11 Free energy change for predicting feasibility of a reaction
P.P. 4.18 to 4.23
4.12 Standard free energy change
4.13 Standard free energy change and equilibrium constant
P.P. 4.24 to 4.31
4.14 Gibbs free energy change and nonmechanical work
4.15 Absolute entropies and third law of thermodynamics
P.P 4.35 to 4.36
Additional numerical problems for practice 1 to 10
Conceptual Questions with Answers: 12
Revision Exercises
Very Short Answer questions 25
Short Answer Questions 48
Long Answer Questions 10
Competition File
Numerical problems for competitive examinations 17
Objective Questions: 36
Fill in the blanks: 10
True or False: 10
Study Plan
Day 1
4.1 Some basic terms and concepts
4.2 Internal energy, heat and work
4.3 First law of thermodynamics
Day 2
Practice Problems 4.1 to 4.6
4.4 Enthalpy and Enthalpy change
P.P 4.7 to 4.10
Day 3
4.5 Limitations of first law of thermodynamics: Need of second law
4.6 Randomness and spontaneity
4.7 Entropy
4.8 Second law of thermodynamics
4.9 Entropy change during phase transition
Day 4
P.P. 4.11 to 4.17
4.10 Gibbs Free energy and free energy change
4.11 Free energy change for predicting feasibility of a reaction
Day 5
P.P. 4.18 to 4.23
4.12 Standard free energy change
4.13 Standard free energy change and equilibrium constant
P.P. 4.24 to 4.31
Day 6
4.14 Gibbs free energy change and nonmechanical work
4.15 Absolute entropies and third law of thermodynamics
P.P 4.35 to 4.36
Day 7
Additional numerical problems for practice 1 to 10
Conceptual Questions with Answers: 12
Day 8
Revision Exercises: Very Short Answer questions 25
Day 9
Revision Exercises: Short Answer Questions 1 t o 24
Day 10
Revision Exercises: Short Answer Questions 25 to 48
Day 11
Competition File: Numerical problems for competitive examinations 17
Day 12
Competition File: Objective Questions: 1 to 18
Day 13
Competition File: Objective Questions: 19to 36
Day 14
Competition File: Fill in the blanks: 10
Competition File: True or False: 10
Day 15
Concept Revision
Fromula Revision
Day 16 to 30
Revision of the chapter/Extra problems from other books like R.C. Mukherjee, Test Paper books etc..
Use all available time productively to improve your understanding, ability to solve problems and recollection of principles.
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Updated 3 Feb 2016, 11 March 2009
Showing posts with label Chemical energetics. Show all posts
Showing posts with label Chemical energetics. Show all posts
Wednesday, February 3, 2016
Thursday, May 21, 2015
JEE Main - Core Revision Points - 5. First Law of Thermodynamics and Chemical Energetics
Importance of Core Revision Points: Core Revision Points are important because if you remember them strongly, many more points related to them will come out of your memory and help you to answer question and problems. Read them many times and make sure you remember them very strongly.
JEE Syllabus
Energetics:
First law of thermodynamics;
Internal energy, work and heat,
pressure-volume work;
Enthalpy,
Hess's law;
Heat of reaction, fusion and vapourization;
Second law of thermodynamics;
Entropy;
Free energy;
Criterion of spontaneity.
------------------
Sections in the Chapter - Jauhar
5.1 Some Basic Terms and Concepts
5.2 Modes of Transference of Energy between System and Surroundings
5.3 Internal Energy and Internal Energy Change
5.4 Zeroth Law of Thermodynamics
5.5 Law of Conservation of Energy: First Law of Thermodynamics
5.6 Enthalpy and Enthalpy change
5.7 Exothermic and Endothermic Reactions
5.8 Heat Capacity
5.9 Measurement of Internal Energy (Delta U) and enthalpy (Delta H) of a Reaction
5.10 Thermochemical Equations
5.11 Enthalpy Changes in Chemical Reactions
5.12 Enthalpy of Formation
5.13 enthalpy of Combustion
5.14 Enthalpy of Neutralization
5.15 Enthalpy of phase Transitions
5.16 Hess’s Law of Constant Heat Summation
5.17 Bond Enthalpy
5.18 Sources of Energy
5.19 Alternative Energy Sources
5.1 Some Basic Terms and Concepts
5.2 Modes of Transference of Energy between System and Surroundings
5.3 Internal Energy and Internal Energy Change
5.4 Zeroth Law of Thermodynamics
5.5 Law of Conservation of Energy: First Law of Thermodynamics
First law of thermodynamics;
Energy cannot be created or destroyed.
U = q + w
Internal energy of matter is equal to kinetic energy and potential energy.
The change in internal energy is equal to heat transferred and work done between the system and the surroundings.
Pressure volume work: If the pressure is constant and the matter expands, the work done is given by p * change in volume. This in termed as pressure volume work.
Enthalpy = U + pv
5.6 Enthalpy and Enthalpy change
5.7 Exothermic and Endothermic Reactions
5.8 Heat Capacity
5.9 Measurement of Internal Energy (Delta U) and enthalpy (Delta H) of a Reaction
5.10 Thermochemical Equations
5.11 Enthalpy Changes in Chemical Reactions
5.12 Enthalpy of Formation
5.13 enthalpy of Combustion
5.14 Enthalpy of Neutralization
5.15 Enthalpy of phase Transitions
5.16 Hess’s Law of Constant Heat Summation
Hess's Law
Hess's Law states that the enthalpy change for a reaction that occurs in many steps is the same as if it occurred in one step. Another way to put this is if several reactions add up to some total reaction, then their enthalpy changes will add up to the enthalpy change for the total reaction.
5.17 Bond Enthalpy
5.18 Sources of Energy
5.19 Alternative Energy Sources
First law of thermodynamics;
Energy cannot be created or destroyed.
U = q + w
Internal energy of matter is equal to kinetic energy and potential energy.
The change in internal energy is equal to heat transferred and work done between the system and the surroundings.
Pressure volume work: If the pressure is constant and the matter expands, the work done is given by p * change in volume. This in termed as pressure volume work.
Enthalpy = U + pv
Hess's Law
Hess's Law states that the enthalpy change for a reaction that occurs in many steps is the same as if it occurred in one step. Another way to put this is if several reactions add up to some total reaction, then their enthalpy changes will add up to the enthalpy change for the total reaction.
JEE Syllabus
Energetics:
First law of thermodynamics;
Internal energy, work and heat,
pressure-volume work;
Enthalpy,
Hess's law;
Heat of reaction, fusion and vapourization;
Second law of thermodynamics;
Entropy;
Free energy;
Criterion of spontaneity.
------------------
Sections in the Chapter - Jauhar
5.1 Some Basic Terms and Concepts
5.2 Modes of Transference of Energy between System and Surroundings
5.3 Internal Energy and Internal Energy Change
5.4 Zeroth Law of Thermodynamics
5.5 Law of Conservation of Energy: First Law of Thermodynamics
5.6 Enthalpy and Enthalpy change
5.7 Exothermic and Endothermic Reactions
5.8 Heat Capacity
5.9 Measurement of Internal Energy (Delta U) and enthalpy (Delta H) of a Reaction
5.10 Thermochemical Equations
5.11 Enthalpy Changes in Chemical Reactions
5.12 Enthalpy of Formation
5.13 enthalpy of Combustion
5.14 Enthalpy of Neutralization
5.15 Enthalpy of phase Transitions
5.16 Hess’s Law of Constant Heat Summation
5.17 Bond Enthalpy
5.18 Sources of Energy
5.19 Alternative Energy Sources
5.1 Some Basic Terms and Concepts
5.2 Modes of Transference of Energy between System and Surroundings
5.3 Internal Energy and Internal Energy Change
5.4 Zeroth Law of Thermodynamics
5.5 Law of Conservation of Energy: First Law of Thermodynamics
First law of thermodynamics;
Energy cannot be created or destroyed.
U = q + w
Internal energy of matter is equal to kinetic energy and potential energy.
The change in internal energy is equal to heat transferred and work done between the system and the surroundings.
Pressure volume work: If the pressure is constant and the matter expands, the work done is given by p * change in volume. This in termed as pressure volume work.
Enthalpy = U + pv
5.6 Enthalpy and Enthalpy change
5.7 Exothermic and Endothermic Reactions
5.8 Heat Capacity
5.9 Measurement of Internal Energy (Delta U) and enthalpy (Delta H) of a Reaction
5.10 Thermochemical Equations
5.11 Enthalpy Changes in Chemical Reactions
5.12 Enthalpy of Formation
5.13 enthalpy of Combustion
5.14 Enthalpy of Neutralization
5.15 Enthalpy of phase Transitions
5.16 Hess’s Law of Constant Heat Summation
Hess's Law
Hess's Law states that the enthalpy change for a reaction that occurs in many steps is the same as if it occurred in one step. Another way to put this is if several reactions add up to some total reaction, then their enthalpy changes will add up to the enthalpy change for the total reaction.
5.17 Bond Enthalpy
5.18 Sources of Energy
5.19 Alternative Energy Sources
First law of thermodynamics;
Energy cannot be created or destroyed.
U = q + w
Internal energy of matter is equal to kinetic energy and potential energy.
The change in internal energy is equal to heat transferred and work done between the system and the surroundings.
Pressure volume work: If the pressure is constant and the matter expands, the work done is given by p * change in volume. This in termed as pressure volume work.
Enthalpy = U + pv
Hess's Law
Hess's Law states that the enthalpy change for a reaction that occurs in many steps is the same as if it occurred in one step. Another way to put this is if several reactions add up to some total reaction, then their enthalpy changes will add up to the enthalpy change for the total reaction.
Wednesday, March 11, 2009
IIT JEE - Study Guide - 5. First Law of Thermodynamics and Chemical Energetics
Sections in the Chapter - Jauhar
5.1 Some Basic Terms and Concepts
5.2 Modes of Transference of Energy between System and Surroundings
5.3 Internal Energy and Internal Energy Change
5.4 Zeroth Law of Thermodynamics
5.5 Law of Conservation of Energy: First Law of Thermodynamics
5.6 Enthalpy and Enthalpy change
5.7 Exothermic and Endothermic Reactions
5.8 Heat Capacity
5.9 Measurement of Internal Energy (Delta U) and enthalpy (Delta H) of a Reaction
5.10 Thermochemical Equations
5.11 Enthalpy Changes in Chemical Reactions
5.12 Enthalpy of Formation
5.13 enthalpy of Combustion
5.14 Enthalpy of Neutralization
5.15 Enthalpy of phase Transitions
5.16 Hess’s Law of Constant Heat Summation
5.17 Bond Enthalpy
5.18 Sources of Energy
5.19 Alternative Energy Sources
Conceptual Questions with Answers: 21
Additional Numerical Problems for Practice: 11
Revision Exercises
Very Short Answer questions: 20
Short Answer Questions: 32
Long Answer Questions: 6
Competition File
Numerical Problems 15
Objective Questions: 39
Fill in the blanks: 10
True or False: 10
Study Plan
Day 1
5.1 Some Basic Terms and Concepts
Ex 5.1
5.2 Modes of Transference of Energy Between System and Surroundings
Day 2
5.3 Internal Energy and Internal Energy Change
5.4 Zeroth Law of Thermodynamics
5.5 Law of Conservation of Energy: First Law of Thermodynamics
Day 3
5.6 Enthalpy and Enthalpy change
Ex 5.2 to 5.6
Day 4
Practice Problems 5.1 to 5.5
5.7 Exothermic and Endothermic Reactions
P.P 5.6, 5.7
Day 5
5.8 Heat Capacity
Ex. 5.7 to 5.9
5.9 Measurement of Internal Energy (Delta U) and enthalpy (Delta H) of a Reaction
Day 6
5.10 Thermochemical Equations
5.11 Enthalpy Changes in Chemical Reactions
5.12 Enthalpy of Formation
Day 7
Ex. 5.11 to 5.14
P.P 5.8 to 5.14
Day 8
5.13 enthalpy of Combustion
Ex. 5.15 to 17
P.P 5.15 to 5.18
Day 9
5.14 Enthalpy of Neutralization
Ex. 5.18
5.15 Enthalpy of phase Transitions
Ex. 5.19, 5.20
Day 10
5.16 Hess’s Law of Constant Heat Summation
Ex. 5.22, 5.23
P.P 5.19 to 5.22
Day 11
Hess’s Law continued
P.P 5.23,
Ex. 5.25
P.P 5.24
Ex. 5.26 to 5.28
P.P. 5.25 to 5.29
Day 12
5.17 Bond Enthalpy
Ex. 5.29 to 5.32
Day 13
P.P 5.30 to 5.33
5.18 Sources of Energy
5.19 Alternative Energy Sources
Day 14
Conceptual Questions with Answers: 1 to 21
Day 15
Additional Numerical Problems for Practice: 5.1 to 5.11
Revision Period Starting
Day 16
Revision Exercises: Very Short Answer questions: 20
Day 17
Revision Exercises: Short Answer Questions: 1 to 16
Day 18
Revision Exercises: Short Answer Questions: 16 to 32
Day 19
Competition File: Numerical Problems 1 to 5
Day 20
Competition File: Numerical Problems 6 to 10
Day 21
Competition File: Numerical Problems 11 to 15
Day 22
Objective Questions: 1 to 20
Day 23
Objective Questions: 21 to 39
Day 24
Fill in the blanks: 10
Day 25
True or False: 10
Day 26
Formula Revision
Day 27
Concept Revision
5.1 Some Basic Terms and Concepts
5.2 Modes of Transference of Energy between System and Surroundings
5.3 Internal Energy and Internal Energy Change
5.4 Zeroth Law of Thermodynamics
5.5 Law of Conservation of Energy: First Law of Thermodynamics
5.6 Enthalpy and Enthalpy change
5.7 Exothermic and Endothermic Reactions
5.8 Heat Capacity
5.9 Measurement of Internal Energy (Delta U) and enthalpy (Delta H) of a Reaction
5.10 Thermochemical Equations
5.11 Enthalpy Changes in Chemical Reactions
5.12 Enthalpy of Formation
5.13 enthalpy of Combustion
5.14 Enthalpy of Neutralization
5.15 Enthalpy of phase Transitions
5.16 Hess’s Law of Constant Heat Summation
5.17 Bond Enthalpy
5.18 Sources of Energy
5.19 Alternative Energy Sources
Conceptual Questions with Answers: 21
Additional Numerical Problems for Practice: 11
Revision Exercises
Very Short Answer questions: 20
Short Answer Questions: 32
Long Answer Questions: 6
Competition File
Numerical Problems 15
Objective Questions: 39
Fill in the blanks: 10
True or False: 10
Study Plan
Day 1
5.1 Some Basic Terms and Concepts
Ex 5.1
5.2 Modes of Transference of Energy Between System and Surroundings
Day 2
5.3 Internal Energy and Internal Energy Change
5.4 Zeroth Law of Thermodynamics
5.5 Law of Conservation of Energy: First Law of Thermodynamics
Day 3
5.6 Enthalpy and Enthalpy change
Ex 5.2 to 5.6
Day 4
Practice Problems 5.1 to 5.5
5.7 Exothermic and Endothermic Reactions
P.P 5.6, 5.7
Day 5
5.8 Heat Capacity
Ex. 5.7 to 5.9
5.9 Measurement of Internal Energy (Delta U) and enthalpy (Delta H) of a Reaction
Day 6
5.10 Thermochemical Equations
5.11 Enthalpy Changes in Chemical Reactions
5.12 Enthalpy of Formation
Day 7
Ex. 5.11 to 5.14
P.P 5.8 to 5.14
Day 8
5.13 enthalpy of Combustion
Ex. 5.15 to 17
P.P 5.15 to 5.18
Day 9
5.14 Enthalpy of Neutralization
Ex. 5.18
5.15 Enthalpy of phase Transitions
Ex. 5.19, 5.20
Day 10
5.16 Hess’s Law of Constant Heat Summation
Ex. 5.22, 5.23
P.P 5.19 to 5.22
Day 11
Hess’s Law continued
P.P 5.23,
Ex. 5.25
P.P 5.24
Ex. 5.26 to 5.28
P.P. 5.25 to 5.29
Day 12
5.17 Bond Enthalpy
Ex. 5.29 to 5.32
Day 13
P.P 5.30 to 5.33
5.18 Sources of Energy
5.19 Alternative Energy Sources
Day 14
Conceptual Questions with Answers: 1 to 21
Day 15
Additional Numerical Problems for Practice: 5.1 to 5.11
Revision Period Starting
Day 16
Revision Exercises: Very Short Answer questions: 20
Day 17
Revision Exercises: Short Answer Questions: 1 to 16
Day 18
Revision Exercises: Short Answer Questions: 16 to 32
Day 19
Competition File: Numerical Problems 1 to 5
Day 20
Competition File: Numerical Problems 6 to 10
Day 21
Competition File: Numerical Problems 11 to 15
Day 22
Objective Questions: 1 to 20
Day 23
Objective Questions: 21 to 39
Day 24
Fill in the blanks: 10
Day 25
True or False: 10
Day 26
Formula Revision
Day 27
Concept Revision
Sunday, December 28, 2008
Energetics- Study Guide - IIT JEE
First law of thermodynamics; Internal energy, work and heat, pressure-volume work; Enthalpy, Hess's law; Heat of reaction, fusion and vapourization; Second law of thermodynamics; Entropy; Free energy; Criterion of spontaneity.
Tuesday, August 12, 2008
Criterion of Spontaneity and Free Energy
July-December Revision
The flow of heat takes from a body at high temperature to a body at low temperature through conduction, convection or radiation. Similarly a liquid at higher level flows to a lower level. In both these cases the action occurs without any additional support. But if a liquid at lower level has to go higher level additional supporting activity is required.
Similarly in chemical reactions some reactions take place if the reactants are in contact. Some reactions will not take place through contact but require additional inputs like heat, catalysts etc. Reactions that take place due to contact alone are called spontaneous reactions. The rate of reaction is not the issue here. Even if the rate of reaction is very slow, if the reaction is taking place, it is a spontaneous reaction.
What determines the spontaneity of a chemical reaction?
Is decrease in enthalpy in the reaction a criterion for spontaneity?
In exothermic reactions, enthalpy of products is less than that of reactants. Thus some persons postulated that a spontaneous chemical reaction may be due to decrease in energy of the products. It sounds reasonable. But some scientists found that some endothermic reactions are also spontaneous. Therefore it is concluded that enthalpy may be a contributory factor for spontaneity, but it is not the complete explanation.
Is entropy a criterion for spontaneity?
Entropy is a thermodynamic function. It can be interpreted as measuring disorder in the system. A gas is more disordered than a liquid and a liquid is more disordered than a solid. In a chemical reaction, if the disorder in products is more than that of reactants, entropy increases. It is found in examples like diffusion of gases etc. that in spontaneous activities disorder increases.
As heat is added to the system, solids become liquids and liquids become gases. Hence heat increases entropy. Entropy is defined as
ΔS = qrev/T for a reversible reaction.
The criterion of spontaneity is defined by the total entropy change of system and surrounding. The total entropy change (ΔStotal) for the system and surroundings of a spontaneous process is given by
ΔStotal = ΔSsystem + ΔSsurrounding > 0
Gibbs energy of Free energy
Gibbs energy or Gibbs function is a thermodynamic functions defined by
G = H-TS
For a constant temperature reaction
ΔGsys = ΔHsys -T ΔSsys
Criterion for spontaneity in terms of Gibbs energy or function is that
If ΔG is negative or< 0, the reaction will be spontaneous.
This condition comes from the condition that was given above only. That is
ΔStotal = ΔSsystem + ΔSsurrounding > 0
ΔSsurr = ΔHsurr/T = -ΔHsys/T (because what system loses surrounding gains and vice versa)
Hence
ΔStotal = ΔSsystem - ΔHsys/T
=> TΔStotal = TΔSsystem - ΔHsys
As spontaneity criterion is ΔStotal > 0
RHS must be greater than 0.
=> TΔSsystem - ΔHsys > 0
=> -( ΔHsys - TΔSsystem) > 0
=> ( ΔHsys - TΔSsystem) < 0
The flow of heat takes from a body at high temperature to a body at low temperature through conduction, convection or radiation. Similarly a liquid at higher level flows to a lower level. In both these cases the action occurs without any additional support. But if a liquid at lower level has to go higher level additional supporting activity is required.
Similarly in chemical reactions some reactions take place if the reactants are in contact. Some reactions will not take place through contact but require additional inputs like heat, catalysts etc. Reactions that take place due to contact alone are called spontaneous reactions. The rate of reaction is not the issue here. Even if the rate of reaction is very slow, if the reaction is taking place, it is a spontaneous reaction.
What determines the spontaneity of a chemical reaction?
Is decrease in enthalpy in the reaction a criterion for spontaneity?
In exothermic reactions, enthalpy of products is less than that of reactants. Thus some persons postulated that a spontaneous chemical reaction may be due to decrease in energy of the products. It sounds reasonable. But some scientists found that some endothermic reactions are also spontaneous. Therefore it is concluded that enthalpy may be a contributory factor for spontaneity, but it is not the complete explanation.
Is entropy a criterion for spontaneity?
Entropy is a thermodynamic function. It can be interpreted as measuring disorder in the system. A gas is more disordered than a liquid and a liquid is more disordered than a solid. In a chemical reaction, if the disorder in products is more than that of reactants, entropy increases. It is found in examples like diffusion of gases etc. that in spontaneous activities disorder increases.
As heat is added to the system, solids become liquids and liquids become gases. Hence heat increases entropy. Entropy is defined as
ΔS = qrev/T for a reversible reaction.
The criterion of spontaneity is defined by the total entropy change of system and surrounding. The total entropy change (ΔStotal) for the system and surroundings of a spontaneous process is given by
ΔStotal = ΔSsystem + ΔSsurrounding > 0
Gibbs energy of Free energy
Gibbs energy or Gibbs function is a thermodynamic functions defined by
G = H-TS
For a constant temperature reaction
ΔGsys = ΔHsys -T ΔSsys
Criterion for spontaneity in terms of Gibbs energy or function is that
If ΔG is negative or< 0, the reaction will be spontaneous.
This condition comes from the condition that was given above only. That is
ΔStotal = ΔSsystem + ΔSsurrounding > 0
ΔSsurr = ΔHsurr/T = -ΔHsys/T (because what system loses surrounding gains and vice versa)
Hence
ΔStotal = ΔSsystem - ΔHsys/T
=> TΔStotal = TΔSsystem - ΔHsys
As spontaneity criterion is ΔStotal > 0
RHS must be greater than 0.
=> TΔSsystem - ΔHsys > 0
=> -( ΔHsys - TΔSsystem) > 0
=> ( ΔHsys - TΔSsystem) < 0
Sunday, February 3, 2008
IIT JEE Revision Ch. 6 ENERGETICS Core Points
JEE Syllabus
Energetics:
First law of thermodynamics;
Internal energy, work and heat,
pressure-volume work;
Enthalpy,
Hess's law;
Heat of reaction, fusion and vapourization;
Second law of thermodynamics;
Entropy;
Free energy;
Criterion of spontaneity.
------------------
First law of thermodynamics;
Energy cannot be created or destroyed.
U = q + w
Internal energy of matter is equal to kinetic energy and potential energy.
The change in internal energy is equal to heat transferred and work done between the system and the surroundings.
Pressure volume work: If the pressure is constant and the matter expands, the work done is given by p * change in volume. This in termed as pressure volume work.
Enthalpy = U + pv
Hess's Law
Hess's Law states that the enthalpy change for a reaction that occurs in many steps is the same as if it occurred in one step. Another way to put this is if several reactions add up to some total reaction, then their enthalpy changes will add up to the enthalpy change for the total reaction.
Second law of thermodynamics
In general it is impossible to perform a transformation whose only final result is to convert into useful work heat extracted from a source that is at the same temperature throughout. This statement is Lord Kelvin's version of the second law of thermodynamics.
Another version of this law, formulated by R. J. E. Clausius, states that a transformation is impossible whose only final result is to transfer heat from a body at a given temperature to a body at higher temperature; in other words, the spontaneous flow of heat from hot to cold bodies is reversible only with the expenditure of mechanical or other nonthermal energy.
The Second law of Thermodynamics states that every spontaneous change is accompanied by an increase in entropy which is a measure of the randomness or disorder of a system.
For each topic in the JEE syllabus detailed revision points are posted in the following posts.
Energetics:
First law of thermodynamics;
Internal energy, work and heat,
pressure-volume work;
Enthalpy,
Hess's law;
Heat of reaction, fusion and vapourization;
Second law of thermodynamics;
Entropy;
Free energy;
Criterion of spontaneity.
------------------
First law of thermodynamics;
Energy cannot be created or destroyed.
U = q + w
Internal energy of matter is equal to kinetic energy and potential energy.
The change in internal energy is equal to heat transferred and work done between the system and the surroundings.
Pressure volume work: If the pressure is constant and the matter expands, the work done is given by p * change in volume. This in termed as pressure volume work.
Enthalpy = U + pv
Hess's Law
Hess's Law states that the enthalpy change for a reaction that occurs in many steps is the same as if it occurred in one step. Another way to put this is if several reactions add up to some total reaction, then their enthalpy changes will add up to the enthalpy change for the total reaction.
Second law of thermodynamics
In general it is impossible to perform a transformation whose only final result is to convert into useful work heat extracted from a source that is at the same temperature throughout. This statement is Lord Kelvin's version of the second law of thermodynamics.
Another version of this law, formulated by R. J. E. Clausius, states that a transformation is impossible whose only final result is to transfer heat from a body at a given temperature to a body at higher temperature; in other words, the spontaneous flow of heat from hot to cold bodies is reversible only with the expenditure of mechanical or other nonthermal energy.
The Second law of Thermodynamics states that every spontaneous change is accompanied by an increase in entropy which is a measure of the randomness or disorder of a system.
For each topic in the JEE syllabus detailed revision points are posted in the following posts.
JEE Revision - Chemical Energetics - Introduction
Thermodynamics is the study of energy and energy changes.
Energy is the capacity to do work.
Energy is present in variou forms like, potential energy, kinetic energy, electrical etc.
Energy and energy changes are measured in terms of heat energy.
The study of the flow of heat or any other form of energy into or out of a system as it undergoes a physical aor chemical transformation is called thermodynamics.
The branch of thermodynamics dealing with energy changes accompanying chemical transformation is called chemical thermodynamics.
Concepts
System
surroundings
boundary
Open system: A system which can exchange matter as well as energy with the surroundings is called an open system.
Closed system: A system which can exchange energy but not matter with the surroundings is called a closed system.
Isolated system: A system which can neither exchange matter nor energy with the surroundings is called an isolated system.
Processes
Isothermal: Temperature of the system is constant.
Adiabatic: No heat flows into or out of the system.
Isochoric: volume of the system remains the same.
Isobaric: Pressure of the sytem remains the same.
Reversible: The system changes in infinitesimal steps and they can be reversed.
Irrevesible: Real life systems do not satisfy the reverbility criterion and hence irreversible.
Modes of transfer of energy between system and surroundings
1. Heat (Q): Energy is exchanged between the system and the surroundings as heat if they are at different temperatures.
2. Another modes of transfer of energy is work. Work is said to be performed if th point of application of a force is displaced in the direction of the force.
Pressure volume work
Pressure volume work is mechanical work. It is the work done when the gas expands or contracts against external pressure.
It is equal to force multiplied by distance moved or pressured mulitiplied by change in volume.
Energy is the capacity to do work.
Energy is present in variou forms like, potential energy, kinetic energy, electrical etc.
Energy and energy changes are measured in terms of heat energy.
The study of the flow of heat or any other form of energy into or out of a system as it undergoes a physical aor chemical transformation is called thermodynamics.
The branch of thermodynamics dealing with energy changes accompanying chemical transformation is called chemical thermodynamics.
Concepts
System
surroundings
boundary
Open system: A system which can exchange matter as well as energy with the surroundings is called an open system.
Closed system: A system which can exchange energy but not matter with the surroundings is called a closed system.
Isolated system: A system which can neither exchange matter nor energy with the surroundings is called an isolated system.
Processes
Isothermal: Temperature of the system is constant.
Adiabatic: No heat flows into or out of the system.
Isochoric: volume of the system remains the same.
Isobaric: Pressure of the sytem remains the same.
Reversible: The system changes in infinitesimal steps and they can be reversed.
Irrevesible: Real life systems do not satisfy the reverbility criterion and hence irreversible.
Modes of transfer of energy between system and surroundings
1. Heat (Q): Energy is exchanged between the system and the surroundings as heat if they are at different temperatures.
2. Another modes of transfer of energy is work. Work is said to be performed if th point of application of a force is displaced in the direction of the force.
Pressure volume work
Pressure volume work is mechanical work. It is the work done when the gas expands or contracts against external pressure.
It is equal to force multiplied by distance moved or pressured mulitiplied by change in volume.
JEE Revision - First law of thermodynamics
First law of thermodynamics;
Energy cannot be created or destroyed.
Energy can neither be created or nor destroyed although it can be converted from one form into another.
The energy of a system that isolated from its surroundings is constant.
Mathematical expresson for the first law
ΔU = q + w
q = heat added to the sytem
w = work done on the system
Sign conventions for heat and work
When w and q are positive, the internal energy increases. It means that energy is supplied to the system.
When w and q are negative, the internal energy decreases. It means that energy has left the system.
Energy cannot be created or destroyed.
Energy can neither be created or nor destroyed although it can be converted from one form into another.
The energy of a system that isolated from its surroundings is constant.
Mathematical expresson for the first law
ΔU = q + w
q = heat added to the sytem
w = work done on the system
Sign conventions for heat and work
When w and q are positive, the internal energy increases. It means that energy is supplied to the system.
When w and q are negative, the internal energy decreases. It means that energy has left the system.
JEE Revision - Internal energy, work and heat
every substance possesses definite quantity of energy which depends upon factors such a chemical nature of the substance, temperature and pressure. This is known as internal energy or intrinsic energy and is represented by the symbol U.
Internal energy is made up of kinetic energy and potential energy of the constituent particles. The kinetic energy arises due to motion of its particles and includes their translational motion energy, rorational motion energy, and vibrational motion energy etc. We studied various potential energies in physics. The total potential energy of a substance arises from different interactions between the particles due to gravitation, coulomb forces, magentic forces, and nuclear forces.
It is important to note that different substances have different internal energies even at the same temperature and volume. This fact is important when we talk of internal energy of reactants and internal energy of products being different at constant temperature and volume. Internal energy of a substance depends upon the nature of constituting atoms and bonds, apart from temperature and volume.
while internal energy cannot be measured, change in internal energy during a process may identified.
ΔU = Up - Ur
ΔU = change in internal energy of a chemical reaction
Up = internal energy of products
Ur = internal energy of reactants
If internal energy of products is more than the internal energy of reactants, the reaction requires input of energy in the form of heat (endothermic reaction).
If internal energy of products is less e than the internal energy of reactants, the reaction gives out energy in the form of heat (exothermic reaction).
U = q + w
Internal energy of matter is equal to kinetic energy and potential energy.
The change in internal energy is equal to heat transferred and work done between the system and the surroundings.
Internal energy is made up of kinetic energy and potential energy of the constituent particles. The kinetic energy arises due to motion of its particles and includes their translational motion energy, rorational motion energy, and vibrational motion energy etc. We studied various potential energies in physics. The total potential energy of a substance arises from different interactions between the particles due to gravitation, coulomb forces, magentic forces, and nuclear forces.
It is important to note that different substances have different internal energies even at the same temperature and volume. This fact is important when we talk of internal energy of reactants and internal energy of products being different at constant temperature and volume. Internal energy of a substance depends upon the nature of constituting atoms and bonds, apart from temperature and volume.
while internal energy cannot be measured, change in internal energy during a process may identified.
ΔU = Up - Ur
ΔU = change in internal energy of a chemical reaction
Up = internal energy of products
Ur = internal energy of reactants
If internal energy of products is more than the internal energy of reactants, the reaction requires input of energy in the form of heat (endothermic reaction).
If internal energy of products is less e than the internal energy of reactants, the reaction gives out energy in the form of heat (exothermic reaction).
U = q + w
Internal energy of matter is equal to kinetic energy and potential energy.
The change in internal energy is equal to heat transferred and work done between the system and the surroundings.
Revision - Pressure-Volume Work
Heat and Work
How one can measure the energy transferred as heat and work in and out of a system.
Work
A common type of work associated by a chemical process is through gas expansion or compression.
Energy produced from the combustion of gasoline (petrol). Gasoline combustion is used to create expanding gases in the cylinders of your car's engine that push out the pistons. This motion is then translated into the motion of the car.
If the expansion is done constant pressure p∆V gives the work done. This is termed pressure-volume-work.
How one can measure the energy transferred as heat and work in and out of a system.
Work
A common type of work associated by a chemical process is through gas expansion or compression.
Energy produced from the combustion of gasoline (petrol). Gasoline combustion is used to create expanding gases in the cylinders of your car's engine that push out the pistons. This motion is then translated into the motion of the car.
If the expansion is done constant pressure p∆V gives the work done. This is termed pressure-volume-work.
JEE Revision - Enthalpy
For constant temperature and pressure processes, work is done by the system to expand or contract. To study the heat change accompanying chemical reactions at constant pressure and constant temperature, a new term called enthalpy is used.
Enthalpy (H) is a more general measure of the energy of a system. It takes into account the internal energy of a system as well as its pressure and volume.
H = U + PV
Like in the case of internal energy, change in enthalpy is identified.
Change in enthalpy is equal to heat supplied to the system
Therefore changes in enthalpy can be measured by changes in heat energy. Enthalpy is also defined as heat flow.
Heat (q) is a type of energy that is related to temperature but is not the same as temperature, since the energy may be used for something other than increasing temperature. For example, the heat energy required to melt 1 mole of molecules is the heat of fusion (H-fus), but no temperature change is involved in this phase change. The heat energy required to change 1 mole of liquid to 1 mole of gas is the heat of vaporization (H-vap). Similarly, no temperature change is involved in this or any other phase change.
It also requires different amounts of heat energy to change the temperature different types of molecules. The energy required to increase the temperature of 1 mole (n) of substance 1 °C is its heat capacity (cp). Similarly, the energy required to increase the temperature of 1 g of substance 1 °C is specific heat. Therefore the heat (q) required for a temperature change (T) can be calculated from
q = n*(cp)*T where (cp = heat capacity; n = number of moles)
Chemical reactions also involve a change in enthalpy (∆H). The ∆Hrxn is the energy per mole of reaction (q/mol). Exothermic reactions release energy into the surrounding and are designated by a negative sign on the H. Endothermic reactions take energy from the surroundings and are designated by a positive value of H.
The value of H can be determined experimentally by measuring the temperature change of the surroundings. Calorimetry is an experiment to measure q by measuring temperature change. The heat capacity of the calorimeter (the constant pressure surroundings in the experiment) is used to convert temperature change to heat energy. Since the heat capacity (CP) refers to the whole calorimeter, heat can be calculated from
q = (cp)T
The value of ∆H can also be calculated theoretically. Since the energy change in a chemical reaction comes from making and breaking bonds, the value of ∆H can be calculated from the energy of the bonds, bond energy. Energy is released (–∆H) when bonds are made and energy is absorbed (+∆H) when bonds are broken.
The value of ∆H can also be calculated from the ∆H values of other reactions. Because H is a state function, the path does not affect the final result. Thus ∆H values of a series of known reactions can be mixed to obtain the ∆H of an unknown reaction. The ∆H value depends on how the reaction is written. If the reaction is reversed, the sign on ∆H is changed. If the stoichiometric coefficients of a reaction are multiplied by some factor, ∆H is multiplied by the same factor. If reactions are added together, so are the ∆H values. Hess's law states the ∆H of a reaction that is the sum of other reactions is the sum of the ∆H values of those reactions.
Enthalpy (H) is a more general measure of the energy of a system. It takes into account the internal energy of a system as well as its pressure and volume.
H = U + PV
Like in the case of internal energy, change in enthalpy is identified.
Change in enthalpy is equal to heat supplied to the system
Therefore changes in enthalpy can be measured by changes in heat energy. Enthalpy is also defined as heat flow.
Heat (q) is a type of energy that is related to temperature but is not the same as temperature, since the energy may be used for something other than increasing temperature. For example, the heat energy required to melt 1 mole of molecules is the heat of fusion (H-fus), but no temperature change is involved in this phase change. The heat energy required to change 1 mole of liquid to 1 mole of gas is the heat of vaporization (H-vap). Similarly, no temperature change is involved in this or any other phase change.
It also requires different amounts of heat energy to change the temperature different types of molecules. The energy required to increase the temperature of 1 mole (n) of substance 1 °C is its heat capacity (cp). Similarly, the energy required to increase the temperature of 1 g of substance 1 °C is specific heat. Therefore the heat (q) required for a temperature change (T) can be calculated from
q = n*(cp)*T where (cp = heat capacity; n = number of moles)
Chemical reactions also involve a change in enthalpy (∆H). The ∆Hrxn is the energy per mole of reaction (q/mol). Exothermic reactions release energy into the surrounding and are designated by a negative sign on the H. Endothermic reactions take energy from the surroundings and are designated by a positive value of H.
The value of H can be determined experimentally by measuring the temperature change of the surroundings. Calorimetry is an experiment to measure q by measuring temperature change. The heat capacity of the calorimeter (the constant pressure surroundings in the experiment) is used to convert temperature change to heat energy. Since the heat capacity (CP) refers to the whole calorimeter, heat can be calculated from
q = (cp)T
The value of ∆H can also be calculated theoretically. Since the energy change in a chemical reaction comes from making and breaking bonds, the value of ∆H can be calculated from the energy of the bonds, bond energy. Energy is released (–∆H) when bonds are made and energy is absorbed (+∆H) when bonds are broken.
The value of ∆H can also be calculated from the ∆H values of other reactions. Because H is a state function, the path does not affect the final result. Thus ∆H values of a series of known reactions can be mixed to obtain the ∆H of an unknown reaction. The ∆H value depends on how the reaction is written. If the reaction is reversed, the sign on ∆H is changed. If the stoichiometric coefficients of a reaction are multiplied by some factor, ∆H is multiplied by the same factor. If reactions are added together, so are the ∆H values. Hess's law states the ∆H of a reaction that is the sum of other reactions is the sum of the ∆H values of those reactions.
Revision - Hess's law
Hess's Law states that the enthalpy change for a reaction that occurs in many steps is the same as if it occurred in one step.
Another way to put this is if several reactions add up to some total reaction, then their enthalpy changes will add up to the enthalpy change for the total reaction.
∆Htotal = ∆Hrxn 1 + ∆Hrxn 2 + ∆Hrxn 3 + etc.
Hess's law problems usually give you two or three reactions with their enthalpy change information, then ask you to find the enthalpy change for some target reaction.
You must figure out how to make the given reactions add up to the target. This can mean reversing the reactions (and reversing the sign on the enthalpy change), or using them multiple times, or both.
Another way to put this is if several reactions add up to some total reaction, then their enthalpy changes will add up to the enthalpy change for the total reaction.
∆Htotal = ∆Hrxn 1 + ∆Hrxn 2 + ∆Hrxn 3 + etc.
Hess's law problems usually give you two or three reactions with their enthalpy change information, then ask you to find the enthalpy change for some target reaction.
You must figure out how to make the given reactions add up to the target. This can mean reversing the reactions (and reversing the sign on the enthalpy change), or using them multiple times, or both.
Revison - Second law of thermodynamics
Second law of thermodynamics
In general it is impossible to perform a transformation whose only final result is to convert into useful work heat extracted from a source that is at the same temperature throughout. This statement is Lord Kelvin's version of the second law of thermodynamics.
Another version of this law, formulated by R. J. E. Clausius, states that a transformation is impossible whose only final result is to transfer heat from a body at a given temperature to a body at higher temperature; in other words, the spontaneous flow of heat from hot to cold bodies is reversible only with the expenditure of mechanical or other nonthermal energy.
The Second law of Thermodynamics states that every spontaneous change is accompanied by an increase in entropy which is a measure of the randomness or disorder of a system.
In general it is impossible to perform a transformation whose only final result is to convert into useful work heat extracted from a source that is at the same temperature throughout. This statement is Lord Kelvin's version of the second law of thermodynamics.
Another version of this law, formulated by R. J. E. Clausius, states that a transformation is impossible whose only final result is to transfer heat from a body at a given temperature to a body at higher temperature; in other words, the spontaneous flow of heat from hot to cold bodies is reversible only with the expenditure of mechanical or other nonthermal energy.
The Second law of Thermodynamics states that every spontaneous change is accompanied by an increase in entropy which is a measure of the randomness or disorder of a system.
JEE Revision - Entropy
Entropy
Entropy (S) is a measure of the amount of disorder in a substance Gases with their rapid random motion are high in entropy, and solids with their ordered crystalline lattice are low in entropy.
The change in entropy (∆S) is determined just like a heat of formation problem, only use entropy values instead.
∆Srxn = (the sum of ∆Sproducts ) - (the sum of the ∆Sreactants )
Note that the units for entropy are given in J / mol K. ∆S is in J / K, since the moles cancel out.
K stands for Kelvins, the temperature unit on the absolute scale. Also called the Kelvin scale, it is named for Lord Kelvin, who developed it, so the units should be capitalized.
Entropy is temperature affected. It is large at high temperatures, and small at low temperatures. Enthalpy is not temperature affected.
Entropy (S) is a measure of the amount of disorder in a substance Gases with their rapid random motion are high in entropy, and solids with their ordered crystalline lattice are low in entropy.
The change in entropy (∆S) is determined just like a heat of formation problem, only use entropy values instead.
∆Srxn = (the sum of ∆Sproducts ) - (the sum of the ∆Sreactants )
Note that the units for entropy are given in J / mol K. ∆S is in J / K, since the moles cancel out.
K stands for Kelvins, the temperature unit on the absolute scale. Also called the Kelvin scale, it is named for Lord Kelvin, who developed it, so the units should be capitalized.
Entropy is temperature affected. It is large at high temperatures, and small at low temperatures. Enthalpy is not temperature affected.
Revision - Criterion of spontaneity
The criterion of spontaneity has two ideas involved. Decrease in energy and increase in disorder.
Spontaneity means one the chemical reaction is started it will continue further without additional work being done on it.
Both endothermic and exothermic reactions were found to be spontaneous.
Gibbs free energy concept combines both ideas being spontaneity.
Gibbs Free Energy
The second law of thermodynamics provides a criterion for the determination of the direction of a thermodynamic process, depending on the value of ∆S(total).
∆S(total)[entropy change due to a process] requires calculation of the entropy change for the system and the surroundings. Hence it is desirable to have a function of state that predicts the directionality (feasibility) of a process in a system without the need for explicit calculations for the surroundings.
Such a function exists and is known as Gibbs free energy (or Gibbs function) (G), after its originator, John Willard Gibbs.
Derivation of Gibbs free energy
Consider a process being carried out at constant pressure and constant temperature.
Under these conditions,
qp = ∆Hsys for the system,
and for the surroundings, -qp = - ∆Hsys
Now, ∆Ssurr = - ∆Hsys/T
Since ∆S(total) = ∆Ssys + ∆Ssurr,
= (∆Ssys - ∆Hsys)/T
= -(∆Hsys - T∆Ssys)/T
= -∆(Hsys - TSsys)T
(since T is constant)
If we define the term (H - TS) as G, Gibbs free energy, then, ∆S(total) = - ∆G /T (or ∆G = -T∆S(total))
Therefore, ∆Gsys is a criterion of spontaneity or directionality of the system:
Value of ∆G and Feasibility of process
< 0 - Spontaneous
0 - Reversible
> 0 - Non-spontaneous
Gibbs Free Energy Calculations at Constant Temperature
Example 1: Phase transition of ice to water
This example uses ∆G = ∆H - T∆S: Consider the ice – water phase transition
∆Hfus = 6007 J mol-1 Tfus = 273.15 K
Find ∆G at –10°C (263.15 K), 0°C (273.15 K) and 10°C (283.15 K)
For all the calculations here, ∆S = ∆Hfus/Tfus = 6007/273.15
At 263.15 K, ∆G = +6007 - 263.15 x 6007/273.15 = +213 J
At 273.15 K, ∆G = +6007 - 273.15 x 6007/273.15 = 0
At 283.15 K, ∆G = +6007 - 283.15 x 6007/273.15 = -213 J
*Below 0° C, water freezes spontaneously (i.e. the right to left direction).
**At 0° C, ice and water co-exist (they are in equilibrium and the process is reversible).
***Above 0° C, ice melts spontaneously (i.e. the left to right direction).
### It is impossible for water to freeze on its own accord above 0oC, or ice to melt on its own accord below 0°C.
Spontaneity means one the chemical reaction is started it will continue further without additional work being done on it.
Both endothermic and exothermic reactions were found to be spontaneous.
Gibbs free energy concept combines both ideas being spontaneity.
Gibbs Free Energy
The second law of thermodynamics provides a criterion for the determination of the direction of a thermodynamic process, depending on the value of ∆S(total).
∆S(total)[entropy change due to a process] requires calculation of the entropy change for the system and the surroundings. Hence it is desirable to have a function of state that predicts the directionality (feasibility) of a process in a system without the need for explicit calculations for the surroundings.
Such a function exists and is known as Gibbs free energy (or Gibbs function) (G), after its originator, John Willard Gibbs.
Derivation of Gibbs free energy
Consider a process being carried out at constant pressure and constant temperature.
Under these conditions,
qp = ∆Hsys for the system,
and for the surroundings, -qp = - ∆Hsys
Now, ∆Ssurr = - ∆Hsys/T
Since ∆S(total) = ∆Ssys + ∆Ssurr,
= (∆Ssys - ∆Hsys)/T
= -(∆Hsys - T∆Ssys)/T
= -∆(Hsys - TSsys)T
(since T is constant)
If we define the term (H - TS) as G, Gibbs free energy, then, ∆S(total) = - ∆G /T (or ∆G = -T∆S(total))
Therefore, ∆Gsys is a criterion of spontaneity or directionality of the system:
Value of ∆G and Feasibility of process
< 0 - Spontaneous
0 - Reversible
> 0 - Non-spontaneous
Gibbs Free Energy Calculations at Constant Temperature
Example 1: Phase transition of ice to water
This example uses ∆G = ∆H - T∆S: Consider the ice – water phase transition
∆Hfus = 6007 J mol-1 Tfus = 273.15 K
Find ∆G at –10°C (263.15 K), 0°C (273.15 K) and 10°C (283.15 K)
For all the calculations here, ∆S = ∆Hfus/Tfus = 6007/273.15
At 263.15 K, ∆G = +6007 - 263.15 x 6007/273.15 = +213 J
At 273.15 K, ∆G = +6007 - 273.15 x 6007/273.15 = 0
At 283.15 K, ∆G = +6007 - 283.15 x 6007/273.15 = -213 J
*Below 0° C, water freezes spontaneously (i.e. the right to left direction).
**At 0° C, ice and water co-exist (they are in equilibrium and the process is reversible).
***Above 0° C, ice melts spontaneously (i.e. the left to right direction).
### It is impossible for water to freeze on its own accord above 0oC, or ice to melt on its own accord below 0°C.
Saturday, January 19, 2008
IIT JEE Ch. 6 ENERGETICS Core Points for Revision
JEE Syllabus
Energetics:
First law of thermodynamics;
Internal energy, work and heat,
pressure-volume work;
Enthalpy,
Hess's law;
Heat of reaction, fusion and vapourization;
Second law of thermodynamics;
Entropy;
Free energy;
Criterion of spontaneity.
------------------
First law of thermodynamics;
Energy cannot be created or destroyed.
U = q + w
Internal energy of matter is equal to kinetic energy and potential energy.
The change in internal energy is equal to heat transferred and work done between the system and the surroundings.
Pressure volume work: If the pressure is constant and the matter expands, the work done is given by p * change in volume. This in termed as pressure volume work.
Enthalpy = U + pv
Hess's Law
Hess's Law states that the enthalpy change for a reaction that occurs in many steps is the same as if it occurred in one step. Another way to put this is if several reactions add up to some total reaction, then their enthalpy changes will add up to the enthalpy change for the total reaction.
Second law of thermodynamics
In general it is impossible to perform a transformation whose only final result is to convert into useful work heat extracted from a source that is at the same temperature throughout. This statement is Lord Kelvin's version of the second law of thermodynamics. Another version of this law, formulated by R. J. E. Clausius, states that a transformation is impossible whose only final result is to transfer heat from a body at a given temperature to a body at higher temperature; in other words, the spontaneous flow of heat from hot to cold bodies is reversible only with the expenditure of mechanical or other nonthermal energy.
The Second law of Thermodynamics states that every spontaneous change is accompanied by an increase in entropy which is a measure of the randomness or disorder of a system.
Energetics:
First law of thermodynamics;
Internal energy, work and heat,
pressure-volume work;
Enthalpy,
Hess's law;
Heat of reaction, fusion and vapourization;
Second law of thermodynamics;
Entropy;
Free energy;
Criterion of spontaneity.
------------------
First law of thermodynamics;
Energy cannot be created or destroyed.
U = q + w
Internal energy of matter is equal to kinetic energy and potential energy.
The change in internal energy is equal to heat transferred and work done between the system and the surroundings.
Pressure volume work: If the pressure is constant and the matter expands, the work done is given by p * change in volume. This in termed as pressure volume work.
Enthalpy = U + pv
Hess's Law
Hess's Law states that the enthalpy change for a reaction that occurs in many steps is the same as if it occurred in one step. Another way to put this is if several reactions add up to some total reaction, then their enthalpy changes will add up to the enthalpy change for the total reaction.
Second law of thermodynamics
In general it is impossible to perform a transformation whose only final result is to convert into useful work heat extracted from a source that is at the same temperature throughout. This statement is Lord Kelvin's version of the second law of thermodynamics. Another version of this law, formulated by R. J. E. Clausius, states that a transformation is impossible whose only final result is to transfer heat from a body at a given temperature to a body at higher temperature; in other words, the spontaneous flow of heat from hot to cold bodies is reversible only with the expenditure of mechanical or other nonthermal energy.
The Second law of Thermodynamics states that every spontaneous change is accompanied by an increase in entropy which is a measure of the randomness or disorder of a system.
Sunday, October 21, 2007
Study Guide Ch. 6 ENERGETICS
JEE Syllabus
Energetics:
First law of thermodynamics;
Internal energy, work and heat,
pressure-volume work;
Enthalpy,
Hess's law;
Heat of reaction, fusion and vapourization;
Second law of thermodynamics;
Entropy;
Free energy;
Criterion of spontaneity.
------------------
Main Topics in TMH Book Chapter
FIRST LAW THERMODYNAMICS
INTERNAL ENERGY AND ENTHALPY
ENTHALPY CHANGE OF A CHEMICAL EQUATION
MOLAR ENHTALPIES OF FORMATIONS
HESS'S LAW OF CONSTANT HEAT SUMMATION
TYPES OF REACTIONS
RELATION BETWEEN DELTA H AND DELTA U OF A CHEMICAL EQUATION
--------------------------
Enthalpy Change, ∆H, With Temperature and State Change
Enthalpy, represented by the symbol H, is essentially a chemistry term for heat, and a term for total kinetic energy of particle motion in a sample. If the reaction is exothermic, the energy contained in the substances is reduced so ∆H has a negative value. On the other hand, if the reaction is endothermic, the substances absorb energy and ∆H is positive.
Temperature is a measure of the average kinetic energy of the particles in a sample.
The faster the molecules in a sample of water the higher the temperature.
Changing temperature of a sample requires a change in enthalpy
It makes sense that you have to add heat to increase temperature, and as a sample cools it gives off heat. The amount of heat required to change the temperature of a sample depends on two factors:
# Mass. The bigger the sample the more heat needed to change its temperature.
# Specific heat capacity of the substance it is made of. This is the energy needed to change the temperature of a 1 gram sample 1 Celsius and it is different for different materials. Its symbol is Cp.
When you go to the beach on a hot summer day, you may have noticed that dry sand can get really hot, while wet sand stays cooler, though the air over them is obviously the same temperature and they are getting the same sunlight. How about dry pavement versus wet pavement? Metal versus tile?
The difference is the specific heat capacity of the different materials. Substances with low specific heat capacities will get very hot very fast, while those with large specific heat capacities warm up much slower since more energy is required to change their temperature. In fact, liquid water has about the highest specific heat of any common substance, so it tends to "heat up" and 'cool down' very slowly compared to most other substances. Ice and steam have lower heat capacities than liquid water.
Some people who prize their outdoor plants set gallon jugs of water around their plants. How does this help the plants? During the heat of the day, that water absorbs energy from the air around it, leaving the air cooler. During the cold night, that water releases that energy as it slowly cools, warming the air around it. This is known as a heat sink.
To find the energy required to change the temperature of a sample use Changing temperature of a sample requires a change in enthalpy
It makes sense that you have to add heat to increase temperature, and as a sample cools it gives off heat. The amount of heat required to change the temperature of a sample depends on two factors:
? Mass. The bigger the sample the more heat needed to change its temperature.
? Specific heat capacity of the substance it is made of. This is the energy needed to change the temperature of a 1 gram sample 1? Celsius and it is different for different materials. For some reason its symbol is Cp. Sorry.
When you go to the beach on a hot summer day, you may have noticed that dry sand can get really hot, while wet sand stays cooler, though the air over them is obviously the same temperature and they are getting the same sunlight. How about dry pavement versus wet pavement? Metal versus tile?
The difference is the specific heat capacity of the different materials. Substances with low specific heat capacities will get very hot very fast, while those with large specific heat capacities warm up much slower since more energy is required to change their temperature. In fact, liquid water has about the highest specific heat of any common substance, so it tends to "heat up" and 'cool down' very slowly compared to most other substances. Ice and steam have lower heat capacities than liquid water.
Some people who prize their outdoor plants set gallon jugs of water around their plants. How does this help the plants? During the heat of the day, that water absorbs energy from the air around it, leaving the air cooler. During the cold night, that water releases that energy as it slowly cools, warming the air around it. This is known as a heat sink.
In Chicago, which sits on the shores of Lake Michigan, the weather near the lake is warmer in winter and cooler in summer than the weather farther inland. This is known as the lake effect. The lake functions as a giant heat sink.
To find the energy required to change the temperature of a sample use ∆H = (m)(∆T)(Cp)
Stoichiometry with Energy
Enthalpy values can be included as part of balanced chemical equations. In exothermic reactions, ∆H is negative, but energy is listed as a positive value on the product side of the equation. In endothermic reactions, ∆H is positive and energy is required and is listed on the reactant side of the equation.
The number part of the energy value (as opposed to the unit) can act as a coefficient of a mole ratio, just as any other coefficient would. This way you can convert energy information to information about any substance in a balanced equation and vice versa.
2H-2 + O-2 → 2 H-2O + 561.6 kJ
Enthalpy changes of different types have different names.
The enthalpy change when something dissolves is heat of solution (∆H sol) The enthalpy change during a chemical reaction is heat of reaction (∆Hrxn). The enthalpy change during a comustion reaction is heat of combustion (∆Hcomb). The enthalpy change during a reaction in which a compound is formed from its elements is heat of formation (∆Hf).
Heat of Formation is an important concept.
Heat of Formation Problems.
The enthalpy change can be calculated by taking the total enthalpy of the products - the total enthalpy of the reactants. The formula is...
DHrxn = (the sum of ∆Hf products ) - (the sum of the ∆Hf reactants )
To do this, multiply the number of moles, or coefficient from the balanced equation, of each substance by its heat of formation), and add them up for the products, then do the same for the reactants. Then subtract. Heats of formation are given in kJ / mol, but ∆H is in kJ, since the moles cancel out.
Hess's Law
Hess's Law states that the enthalpy change for a reaction that occurs in many steps is the same as if it occurred in one step. Another way to put this is if several reactions add up to some total reaction, then their enthalpy changes will add up to the enthalpy change for the total reaction.
DHtotal = ∆Hrxn 1 + ∆Hrxn 2 + ∆Hrxn 3 + etc.
Hess's law problems usually give you two or three reactions with their enthalpy change information, then ask you to find the enthalpy change for some target reaction. You must figure out how to make the given reactions add up to the target. This can mean reversing the reactions (and reversing the sign on the enthalpy change), or using them multiple times, or both.
Entropy
Entropy (S) is a measure of the amount of disorder in a substance Gases with their rapid random motion are high in entropy, and solids with their ordered crystalline lattice are low in entropy.
The change in entropy (∆S) is determined just like a heat of formation problem, only use entropy values instead.
∆Srxn = (the sum of ∆Sproducts ) - (the sum of the ∆Sreactants )
Note that the units for entropy are given in J / mol K. ∆S is in J / K, since the moles cancel out.
K stands for Kelvins, the temperature unit on the absolute scale. Also called the Kelvin scale, it is named for Lord Kelvin, who developed it, so the units should be capitalized.
Entropy is temperature affected. It is large at high temperatures, and small at low temperatures. Enthalpy is not temperature affected.
How Enthalpy and Entropy Drive Change
If a reaction will occur spontaneously, it will occur so the products are said to be favored. If a reaction will not occur spontaneously, then the reverse reaction will(in a reversible reaction), so the reactants are said to be favored.
Exothermic changes, those with negative enthalpy changes, are favored to occur spontaneously. So if ∆H is negative, a reaction is more likely than if it is not.
On the other hand, changes that involve an increase in disorder, or entropy, are favored to occur spontaneously. So if ∆S is positive, a reaction is more likely than if it is not.
If ∆H is negative and ∆S is positive, a reaction will certainly occur, no matter what the temperature. If ∆H is positive and ∆S is negative, there will not be a reaction at any temperature, since both indicators say it wont.
The trouble is that often the two indicators disagree. When they do, the enthalpy tends to win out at low temperatures, and the entropy (since it is temperature affected) tends to win out at high temperatures.
For example, if the enthalpy and entropy change values are both negative, the enthalpy indicates the reaction will occur, and it will at low temperatures. The entropy indicates that there will be no reaction, and at high temperatures there wont. The reverse is true for positive enthalpy and entropy changes. These reactions are more likely to occur at high, but not at low temperatures.
Gibbs Free Energy (∆G) determines for sure whether a reaction will be favored to occur. It is simply a formula that compares ∆H to ∆S in a special way.
∆G = ∆H - T∆S
Temperature must be in Kelvins. If ∆G has a negative value, the reaction will occur spontaneously. If ∆G has a positive value, it will not occur
One complication in calculating the free energy change is that the enthalpy values are typically given in kilojoules (kJ), while entropy values are given in J / K, so you must convert so that both use Joules, or both use kilojoules. It doesn't matter which.
Sometimes you may be asked to find the temperature above which a reaction will or wont occur. This is the temperature at which ∆G is between negative and positive, or when it equals zero.
∆G = 0 so 0 = ∆H - T∆S
Rearrange the equation to solve for T, and you will find that...
T = ∆H/∆S
Above that temperature entropy change determines whether a reaction occurs, and below that temperature, enthalpy change determines whether a reaction occurs.
∆G = 0 means there will be equilibrium in a reversible reaction.
-----------------------------
JEE Question 2007 paper II
For the process (water becoming steam) H-2O(l)(1 bar, 373 K) --> H-2O(g)(1 bar, 373 K), the correct set of thermodynamic parameters is
(A) ∆G=0,∆S= + ve
(B)∆G=0,∆S= -ve
(C)∆G=+ve,∆S=0
(D)∆G=-ve, ∆S= + ve
Solution: A
The answer is A because, because at 100 degree C, the steam and water mixture is at equilibrium. Hence ΔG = 0(G = H - TS), and ΔS is positive. Why? when liquid becomes gas, there is more disorder. More entropy.
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Energetics:
First law of thermodynamics;
Internal energy, work and heat,
pressure-volume work;
Enthalpy,
Hess's law;
Heat of reaction, fusion and vapourization;
Second law of thermodynamics;
Entropy;
Free energy;
Criterion of spontaneity.
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Main Topics in TMH Book Chapter
FIRST LAW THERMODYNAMICS
INTERNAL ENERGY AND ENTHALPY
ENTHALPY CHANGE OF A CHEMICAL EQUATION
MOLAR ENHTALPIES OF FORMATIONS
HESS'S LAW OF CONSTANT HEAT SUMMATION
TYPES OF REACTIONS
RELATION BETWEEN DELTA H AND DELTA U OF A CHEMICAL EQUATION
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Enthalpy Change, ∆H, With Temperature and State Change
Enthalpy, represented by the symbol H, is essentially a chemistry term for heat, and a term for total kinetic energy of particle motion in a sample. If the reaction is exothermic, the energy contained in the substances is reduced so ∆H has a negative value. On the other hand, if the reaction is endothermic, the substances absorb energy and ∆H is positive.
Temperature is a measure of the average kinetic energy of the particles in a sample.
The faster the molecules in a sample of water the higher the temperature.
Changing temperature of a sample requires a change in enthalpy
It makes sense that you have to add heat to increase temperature, and as a sample cools it gives off heat. The amount of heat required to change the temperature of a sample depends on two factors:
# Mass. The bigger the sample the more heat needed to change its temperature.
# Specific heat capacity of the substance it is made of. This is the energy needed to change the temperature of a 1 gram sample 1 Celsius and it is different for different materials. Its symbol is Cp.
When you go to the beach on a hot summer day, you may have noticed that dry sand can get really hot, while wet sand stays cooler, though the air over them is obviously the same temperature and they are getting the same sunlight. How about dry pavement versus wet pavement? Metal versus tile?
The difference is the specific heat capacity of the different materials. Substances with low specific heat capacities will get very hot very fast, while those with large specific heat capacities warm up much slower since more energy is required to change their temperature. In fact, liquid water has about the highest specific heat of any common substance, so it tends to "heat up" and 'cool down' very slowly compared to most other substances. Ice and steam have lower heat capacities than liquid water.
Some people who prize their outdoor plants set gallon jugs of water around their plants. How does this help the plants? During the heat of the day, that water absorbs energy from the air around it, leaving the air cooler. During the cold night, that water releases that energy as it slowly cools, warming the air around it. This is known as a heat sink.
To find the energy required to change the temperature of a sample use Changing temperature of a sample requires a change in enthalpy
It makes sense that you have to add heat to increase temperature, and as a sample cools it gives off heat. The amount of heat required to change the temperature of a sample depends on two factors:
? Mass. The bigger the sample the more heat needed to change its temperature.
? Specific heat capacity of the substance it is made of. This is the energy needed to change the temperature of a 1 gram sample 1? Celsius and it is different for different materials. For some reason its symbol is Cp. Sorry.
When you go to the beach on a hot summer day, you may have noticed that dry sand can get really hot, while wet sand stays cooler, though the air over them is obviously the same temperature and they are getting the same sunlight. How about dry pavement versus wet pavement? Metal versus tile?
The difference is the specific heat capacity of the different materials. Substances with low specific heat capacities will get very hot very fast, while those with large specific heat capacities warm up much slower since more energy is required to change their temperature. In fact, liquid water has about the highest specific heat of any common substance, so it tends to "heat up" and 'cool down' very slowly compared to most other substances. Ice and steam have lower heat capacities than liquid water.
Some people who prize their outdoor plants set gallon jugs of water around their plants. How does this help the plants? During the heat of the day, that water absorbs energy from the air around it, leaving the air cooler. During the cold night, that water releases that energy as it slowly cools, warming the air around it. This is known as a heat sink.
In Chicago, which sits on the shores of Lake Michigan, the weather near the lake is warmer in winter and cooler in summer than the weather farther inland. This is known as the lake effect. The lake functions as a giant heat sink.
To find the energy required to change the temperature of a sample use ∆H = (m)(∆T)(Cp)
Stoichiometry with Energy
Enthalpy values can be included as part of balanced chemical equations. In exothermic reactions, ∆H is negative, but energy is listed as a positive value on the product side of the equation. In endothermic reactions, ∆H is positive and energy is required and is listed on the reactant side of the equation.
The number part of the energy value (as opposed to the unit) can act as a coefficient of a mole ratio, just as any other coefficient would. This way you can convert energy information to information about any substance in a balanced equation and vice versa.
2H-2 + O-2 → 2 H-2O + 561.6 kJ
Enthalpy changes of different types have different names.
The enthalpy change when something dissolves is heat of solution (∆H sol) The enthalpy change during a chemical reaction is heat of reaction (∆Hrxn). The enthalpy change during a comustion reaction is heat of combustion (∆Hcomb). The enthalpy change during a reaction in which a compound is formed from its elements is heat of formation (∆Hf).
Heat of Formation is an important concept.
Heat of Formation Problems.
The enthalpy change can be calculated by taking the total enthalpy of the products - the total enthalpy of the reactants. The formula is...
DHrxn = (the sum of ∆Hf products ) - (the sum of the ∆Hf reactants )
To do this, multiply the number of moles, or coefficient from the balanced equation, of each substance by its heat of formation), and add them up for the products, then do the same for the reactants. Then subtract. Heats of formation are given in kJ / mol, but ∆H is in kJ, since the moles cancel out.
Hess's Law
Hess's Law states that the enthalpy change for a reaction that occurs in many steps is the same as if it occurred in one step. Another way to put this is if several reactions add up to some total reaction, then their enthalpy changes will add up to the enthalpy change for the total reaction.
DHtotal = ∆Hrxn 1 + ∆Hrxn 2 + ∆Hrxn 3 + etc.
Hess's law problems usually give you two or three reactions with their enthalpy change information, then ask you to find the enthalpy change for some target reaction. You must figure out how to make the given reactions add up to the target. This can mean reversing the reactions (and reversing the sign on the enthalpy change), or using them multiple times, or both.
Entropy
Entropy (S) is a measure of the amount of disorder in a substance Gases with their rapid random motion are high in entropy, and solids with their ordered crystalline lattice are low in entropy.
The change in entropy (∆S) is determined just like a heat of formation problem, only use entropy values instead.
∆Srxn = (the sum of ∆Sproducts ) - (the sum of the ∆Sreactants )
Note that the units for entropy are given in J / mol K. ∆S is in J / K, since the moles cancel out.
K stands for Kelvins, the temperature unit on the absolute scale. Also called the Kelvin scale, it is named for Lord Kelvin, who developed it, so the units should be capitalized.
Entropy is temperature affected. It is large at high temperatures, and small at low temperatures. Enthalpy is not temperature affected.
How Enthalpy and Entropy Drive Change
If a reaction will occur spontaneously, it will occur so the products are said to be favored. If a reaction will not occur spontaneously, then the reverse reaction will(in a reversible reaction), so the reactants are said to be favored.
Exothermic changes, those with negative enthalpy changes, are favored to occur spontaneously. So if ∆H is negative, a reaction is more likely than if it is not.
On the other hand, changes that involve an increase in disorder, or entropy, are favored to occur spontaneously. So if ∆S is positive, a reaction is more likely than if it is not.
If ∆H is negative and ∆S is positive, a reaction will certainly occur, no matter what the temperature. If ∆H is positive and ∆S is negative, there will not be a reaction at any temperature, since both indicators say it wont.
The trouble is that often the two indicators disagree. When they do, the enthalpy tends to win out at low temperatures, and the entropy (since it is temperature affected) tends to win out at high temperatures.
For example, if the enthalpy and entropy change values are both negative, the enthalpy indicates the reaction will occur, and it will at low temperatures. The entropy indicates that there will be no reaction, and at high temperatures there wont. The reverse is true for positive enthalpy and entropy changes. These reactions are more likely to occur at high, but not at low temperatures.
Gibbs Free Energy (∆G) determines for sure whether a reaction will be favored to occur. It is simply a formula that compares ∆H to ∆S in a special way.
∆G = ∆H - T∆S
Temperature must be in Kelvins. If ∆G has a negative value, the reaction will occur spontaneously. If ∆G has a positive value, it will not occur
One complication in calculating the free energy change is that the enthalpy values are typically given in kilojoules (kJ), while entropy values are given in J / K, so you must convert so that both use Joules, or both use kilojoules. It doesn't matter which.
Sometimes you may be asked to find the temperature above which a reaction will or wont occur. This is the temperature at which ∆G is between negative and positive, or when it equals zero.
∆G = 0 so 0 = ∆H - T∆S
Rearrange the equation to solve for T, and you will find that...
T = ∆H/∆S
Above that temperature entropy change determines whether a reaction occurs, and below that temperature, enthalpy change determines whether a reaction occurs.
∆G = 0 means there will be equilibrium in a reversible reaction.
-----------------------------
JEE Question 2007 paper II
For the process (water becoming steam) H-2O(l)(1 bar, 373 K) --> H-2O(g)(1 bar, 373 K), the correct set of thermodynamic parameters is
(A) ∆G=0,∆S= + ve
(B)∆G=0,∆S= -ve
(C)∆G=+ve,∆S=0
(D)∆G=-ve, ∆S= + ve
Solution: A
The answer is A because, because at 100 degree C, the steam and water mixture is at equilibrium. Hence ΔG = 0(G = H - TS), and ΔS is positive. Why? when liquid becomes gas, there is more disorder. More entropy.
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