The study of thermodynamics is the study of systems that are too large to be extrapolated by mechanics alone.
For many generations, thermodynamics was vaguely understood, and many of the results were determined only experimentally.
Some of the results posed great theoretical challenges for physicists, who made many unprofitable attempts to explain the origins of the formulas. With the advent of quantum mechanics approached for clarification of the results.
The mechanics of the related particles are immaterially exceedingly complex, nonetheless. For this reason, statistical physics plays an important role in the basis of thermodynamics. Instead of worrying about the exact values of parcels for each particle in a system, we look at statistically averaged values over quantum probabilities.
In fact, introductory concepts like the energy of a system are extrapolated as parameters. The conception of enthalpy is consequential for the temperature and pressure necessitated for any chemical reaction.
Table of Contents
What is Enthalpy?
Heat absorbed or evolved at constant pressure is called heat of the reaction or enthalpy of the reaction. The change in enthalpy accompanying a reaction is called the reaction enthalpy,
Reactant → Product
Enthalpy can also be defined as the total heat content of the system. It’s equal to the sum of internal energy and pressure-volume work. The internal energy change follows the first law of thermodynamics, according to which:
q = ΔU + PΔV
where,
U is internal energy and
H is a static function
Still, ΔV = 0, If the procedure is conveyed out at constant volume. The equation mentioned before additionally reduces to the form:
qv = ΔU
where the subscript v indicates a constant volume.
Therefore, the internal energy change is a constant amount of soaked or expanded heat.
It can be further stated that since ΔU is a state function, thus, qv is also a state function.
If a process is carried out at constant pressure (as is usually the case, because most of the reactions are studied in vessels open to the atmosphere or if a system consists of a gas confined in a cylinder fitted with a piston, the external pressure acting on the piston is the atmospheric pressure), the work of expansion is given by,
w = – PΔV
where ΔV is the increase in volume and P is the constant pressure.
According to the first law of thermodynamics,
q = ΔU – w
where q is the heat soaked by the system, ΔU is the increment in internal energy of the system and w is the work befitted by the system.
Under keeping of constant pressure, emplacing w = -PΔV and depicting the heat absorbed by ap (subscript p denoting constant pressure),
qp =ΔU + PΔV ….(1)
Suppose when the system absorbs q, joules of heat, its internal energy increases from U1 to U2 and the volume increases from V₁ to V₂. Then, we have
ΔU = U2 – U1
and ΔV = V₂ – V₁
Putting these values in equation (1) above, we get
qp = (U₂-U₁) + P(V₂-V₁)
qp = (U₂ + PV₂) – (U₁ + PV₁)
Now, as U, P, and V are the functions of the state, therefore, the quantity U + PV must also be a state function. The thermodynamic quantity U+ PV is called the heat content or enthalpy of the system and is represented by the symbol H, Le, the enthalpy may be defined mathematically by the equation,
H = U + PV ….(2)
Thus, if H2 is the enthalpy of the system in and putting these values in equation (2), we get
H₂ = U₂ + PV₂
H₁ = U₁ + PV₁
qp = H₂ – H1
or
q = ΔH ….(3)
where ΔH = H₂-H1 is the enthalpy change of the system.
Hence, the Enthalpy change of a system is equal to the heat absorbed or evolved by the system at constant pressure. It may be remembered that as most of the reactions are carried out at constant pressure (i.e., in the open vessels), the measured value of the heat evolved or absorbed is the enthalpy change.
Further, putting the value of q, from equation (3) in equation (1), we get
ΔH = ΔU+ PΔV
Enthalpy Change of a Reaction
The enthalpy change can be represented as a function of the increase in internal energy of the system and the pressure-volume work done, i.e., expansion, correspondingly as a process.
The physical concept of enthalpy or heat content – Enthalpy is defined by the subtle expression, H = U + PV, let us try to understand what this volume is. It has been shown earlier that every substance or system has a certain amount of energy stored in it. called internal energy. This energy can be of many types.
The energy stored within the substance or the system that is available for conversion into heat is called the heat content or enthalpy of the substance or the system. Like the internal energy, the absolute value of the heat content or enthalpy of a substance or system cannot be measured and, fortunately, it is not equally necessary. In thermodynamic processes, we are only concerned with enthalpy changes (ΔH), which can be fluently increased experimentally. Subsequently, it can be advertised that U and V are wide parcels, thus the enthalpy is correspondingly a broad property.
Extensive property- Value depends on the volume or size of matter in the system. Exemplifications- mass, volume, internal energy, heat capacity, etc.
Intensive property- Value doesn’t depend on the volume or size of matter in the system. Exemplifications-temperature, viscosity, pressure, etc.
Relationship between Heat of Reaction at Constant Pressure and Constant Volume
It has already been discussed that,
qp = ΔH
and
qv = ΔU ….(1)
It has also been derived already that at constant pressure,
ΔH = AU + PΔV ….(2)
where ΔV is the change in volume
Equation. (2) can be rewritten as:
ΔH = ΔU + P(V₂-V₁) = ΔU + (PV₂-PV₁) ….(3)
where V1 is the initial volume and V₂ is the final volume of the system.
But for ideal gases, PV=nRT so that,
PV₁ = n1 RT
and
PV₂=n₂ RT
where n1 is the number of moles of the gaseous reactants and n2 is the number of moles of the gaseous products.
Substituting these values in the equation. (3),
ΔH = ΔU+ (n2RT- n1RT) = ΔU+(n₂-n1) RT
ΔH = ΔU + Δn RT ….(4)
where Δng = n2 – n1 is the difference between the number of moles of the gaseous products and those of the gaseous reactants. Putting the values of AU from equation (1), equation (4) becomes:
qp= qv+Δn RT
Conditions under which qp=qv or ΔH = ΔU
When the reaction is carried out in a closed vessel so that the volume remains constant, that is, ΔV=0
When the reaction involves only solids or liquids or solutions but no gaseous reactants or products. This is because the change in the volume of solids and liquids during a chemical reaction is negligible.
When reaction involves gaseous reactants and products but their number of moles are equal (ie, = np = nr), e.g., in the reactions
H₂ (g) + Cl₂ (g)→ 2HCl(g)
C(s) + O₂(g) → CO₂ (g)
Thus, qp is different from qv only in those reactions which involve gaseous reactants and products and (np)gaseous ≠ (nr) gaseous.
Sample Problems
Problem 1: The heat of combustion of benzene in a bomb calorimeter (ie constant volume) at 25°C was found to be 3263.9 kJ mol-1. Calculate the heat of combustion of benzene at constant pressure.
Solution:
The reaction is: C6H6(l) + 7 ½O₂ (g) → 6CO₂ (s) + 3H₂O(l)
In this reaction, O2 is the only gaseous reactant and CO2 is the only gaseous product.
Problem 2: If water vapour is assumed to be a complete gas, the molar enthalpy change at 1 bar and 100°C is 41 kJ mol. Calculate the internal energy change when:
1 mole of water vaporizes at 1 bar pressure and 100 °C.
1 mol of water turns into ice.
Solution:
For vaporization of water, the change is: H₂O (l)→ H₂O(g)
For conversion of water into ice, the change is H₂O (l) → H₂O (s) this case, the volume change is negligible.
Hence, ΔH – ΔU = 41.00 kJ mol-1.
Problem 3: A swimmer coming out of a pool is covered with a film of internal energy of 18 g. How much heat must be supplied to make this water evaporate at 298 K? Calculate the evaporation at 100°C. Δvap H0 for water 373 K = 40.66 kJ mol–¹.
Solution:
The process of evaporation is: 18 g H₂O (l) 18 g H₂O (g)
No. of moles in 18 g H₂O = 18 g / 18 g mol-1 = 1 mol
Problem 4: Find the internal energy change for the reaction A(l) → A(g) at 373 K. The heat of vaporization is 40.66 kJ/mol and R=8.3 J mol K–¹.
Solution:
A(l)→ A (g),
ΔH – ΔU + Δng RT
ΔU=ΔH- Δng RT
= 40660 J-1 mol x 8.314J K–¹ mol–¹ x 373 K
= 40660 J-3101 J
= 37559 3 mol
= 37.56 kJ mol–¹
Neeraj Anand, Param Anand
Er. Neeraj K.Anand is a freelance mentor and writer who specializes in Engineering & Science subjects. Neeraj Anand received a B.Tech degree in Electronics and Communication Engineering from N.I.T Warangal & M.Tech Post Graduation from IETE, New Delhi. He has over 30 years of teaching experience and serves as the Head of Department of ANAND CLASSES. He concentrated all his energy and experiences in academics and subsequently grew up as one of the best mentors in the country for students aspiring for success in competitive examinations.
In parallel, he started a Technical Publication "ANAND TECHNICAL PUBLISHERS" in 2002 and Educational Newspaper "NATIONAL EDUCATION NEWS" in 2014 at Jalandhar. Now he is a Director of leading publication "ANAND TECHNICAL PUBLISHERS", "ANAND CLASSES" and "NATIONAL EDUCATION NEWS".
He has published more than hundred books in the field of Physics, Mathematics, Computers and Information Technology. Besides this he has written many books to help students prepare for IIT-JEE and AIPMT entrance exams. He is an executive member of the IEEE (Institute of Electrical & Electronics Engineers. USA) and honorary member of many Indian scientific societies such as Institution of Electronics & Telecommunication Engineers, Aeronautical Society of India, Bioinformatics Institute of India, Institution of Engineers. He has got award from American Biographical Institute Board of International Research in the year 2005.
CBSE Class 11 Chemistry Syllabus is a vast which needs a clear understanding of the concepts and topics. Knowing CBSE Class 11 Chemistry syllabus helps students to understand the course structure of Chemistry.
Unit-wise CBSE Class 11 Syllabus for Chemistry
Below is a list of detailed information on each unit for Class 11 Students.
UNIT I – Some Basic Concepts of Chemistry
General Introduction: Importance and scope of Chemistry.
Nature of matter, laws of chemical combination, Dalton’s atomic theory: concept of elements, atoms and molecules.
Atomic and molecular masses, mole concept and molar mass, percentage composition, empirical and molecular formula, chemical reactions, stoichiometry and calculations based on stoichiometry.
UNIT II – Structure of Atom
Discovery of Electron, Proton and Neutron, atomic number, isotopes and isobars. Thomson’s model and its limitations. Rutherford’s model and its limitations, Bohr’s model and its limitations, concept of shells and subshells, dual nature of matter and light, de Broglie’s relationship, Heisenberg uncertainty principle, concept of orbitals, quantum numbers, shapes of s, p and d orbitals, rules for filling electrons in orbitals – Aufbau principle, Pauli’s exclusion principle and Hund’s rule, electronic configuration of atoms, stability of half-filled and completely filled orbitals.
UNIT III – Classification of Elements and Periodicity in Properties
Significance of classification, brief history of the development of periodic table, modern periodic law and the present form of periodic table, periodic trends in properties of elements -atomic radii, ionic radii, inert gas radii, Ionization enthalpy, electron gain enthalpy, electronegativity, valency. Nomenclature of elements with atomic number greater than 100.
UNIT IV – Chemical Bonding and Molecular Structure
Valence electrons, ionic bond, covalent bond, bond parameters, Lewis structure, polar character of covalent bond, covalent character of ionic bond, valence bond theory, resonance, geometry of covalent molecules, VSEPR theory, concept of hybridization, involving s, p and d orbitals and shapes of some simple molecules, molecular orbital theory of homonuclear diatomic molecules(qualitative idea only), Hydrogen bond.
UNIT V – Chemical Thermodynamics
Concepts of System and types of systems, surroundings, work, heat, energy, extensive and intensive properties, state functions. First law of thermodynamics – internal energy and enthalpy, measurement of U and H, Hess’s law of constant heat summation, enthalpy of bond dissociation, combustion, formation, atomization, sublimation, phase transition, ionization, solution and dilution. Second law of Thermodynamics (brief introduction) Introduction of entropy as a state function, Gibb’s energy change for spontaneous and nonspontaneous processes. Third law of thermodynamics (brief introduction).
UNIT VI – Equilibrium
Equilibrium in physical and chemical processes, dynamic nature of equilibrium, law of mass action, equilibrium constant, factors affecting equilibrium – Le Chatelier’s principle, ionic equilibrium- ionization of acids and bases, strong and weak electrolytes, degree of ionization, ionization of poly basic acids, acid strength, concept of pH, hydrolysis of salts (elementary idea), buffer solution, Henderson Equation, solubility product, common ion effect (with illustrative examples).
UNIT VII – Redox Reactions
Concept of oxidation and reduction, redox reactions, oxidation number, balancing redox reactions, in terms of loss and gain of electrons and change in oxidation number, applications of redox reactions.
UNIT VIII – Organic Chemistry: Some basic Principles and Techniques
General introduction, classification and IUPAC nomenclature of organic compounds. Electronic displacements in a covalent bond: inductive effect, electromeric effect, resonance and hyper conjugation. Homolytic and heterolytic fission of a covalent bond: free radicals, carbocations, carbanions, electrophiles and nucleophiles, types of organic reactions.
UNIT IX – Hydrocarbons
Classification of Hydrocarbons Aliphatic Hydrocarbons: Alkanes – Nomenclature, isomerism, conformation (ethane only), physical properties, chemical reactions. Alkenes – Nomenclature, structure of double bond (ethene), geometrical isomerism, physical properties, methods of preparation, chemical reactions: addition of hydrogen, halogen, water, hydrogen halides (Markovnikov’s addition and peroxide effect), ozonolysis, oxidation, mechanism of electrophilic addition. Alkynes – Nomenclature, structure of triple bond (ethyne), physical properties, methods of preparation, chemical reactions: acidic character of alkynes, addition reaction of – hydrogen, halogens, hydrogen halides and water.
Aromatic Hydrocarbons:
Introduction, IUPAC nomenclature, benzene: resonance, aromaticity, chemical properties: mechanism of electrophilic substitution. Nitration, sulphonation, halogenation, Friedel Craft’s alkylation and acylation, directive influence of functional group in monosubstituted benzene. Carcinogenicity and toxicity.
To know the CBSE Syllabus for all the classes from 1 to 12, visit the Syllabus page of CBSE. Meanwhile, to get the Practical Syllabus of Class 11 Chemistry, read on to find out more about the syllabus and related information in this page.
CBSE Class 11 Chemistry Practical Syllabus with Marking Scheme
In Chemistry subject, practical also plays a vital role in improving their academic scores in the subject. The overall weightage of Chemistry practical mentioned in the CBSE Class 11 Chemistry syllabus is 30 marks. So, students must try their best to score well in practicals along with theory. It will help in increasing their overall academic score.
CBSE Class 11 Chemistry Practical Syllabus
The experiments will be conducted under the supervision of subject teacher. CBSE Chemistry Practicals is for 30 marks. This contribute to the overall practical marks for the subject.
The table below consists of evaluation scheme of practical exams.
Evaluation Scheme
Marks
Volumetric Analysis
08
Salt Analysis
08
Content Based Experiment
06
Project Work
04
Class record and viva
04
Total
30
CBSE Syllabus for Class 11 Chemistry Practical
Micro-chemical methods are available for several of the practical experiments. Wherever possible such techniques should be used.
A. Basic Laboratory Techniques 1. Cutting glass tube and glass rod 2. Bending a glass tube 3. Drawing out a glass jet 4. Boring a cork
B. Characterization and Purification of Chemical Substances 1. Determination of melting point of an organic compound. 2. Determination of boiling point of an organic compound. 3. Crystallization of impure sample of any one of the following: Alum, Copper Sulphate, Benzoic Acid.
C. Experiments based on pH
1. Any one of the following experiments:
Determination of pH of some solutions obtained from fruit juices, solution of known and varied concentrations of acids, bases and salts using pH paper or universal indicator.
Comparing the pH of solutions of strong and weak acids of same concentration.
Study the pH change in the titration of a strong base using universal indicator.
2. Study the pH change by common-ion in case of weak acids and weak bases.
D. Chemical Equilibrium One of the following experiments:
1. Study the shift in equilibrium between ferric ions and thiocyanate ions by increasing/decreasing the concentration of either of the ions. 2. Study the shift in equilibrium between [Co(H2O)6] 2+ and chloride ions by changing the concentration of either of the ions.
E. Quantitative Estimation i. Using a mechanical balance/electronic balance. ii. Preparation of standard solution of Oxalic acid. iii. Determination of strength of a given solution of Sodium hydroxide by titrating it against standard solution of Oxalic acid. iv. Preparation of standard solution of Sodium carbonate. v. Determination of strength of a given solution of hydrochloric acid by titrating it against standard Sodium Carbonatesolution.
F. Qualitative Analysis 1) Determination of one anion and one cation in a given salt Cations‐ Pb2+, Cu2+, As3+, Al3+, Fe3+, Mn2+, Ni2+, Zn2+, Co2+, Ca2+, Sr2+, Ba2+, Mg2+, NH4+ Anions – (CO3)2‐ , S2‐, NO2‐ , SO32‐, SO2‐ , NO ‐ , Cl‐ , Br‐, I‐, PO43‐ , C2O2‐ ,CH3COO‐ (Note: Insoluble salts excluded)
2) Detection of ‐ Nitrogen, Sulphur, Chlorine in organic compounds.
G) PROJECTS Scientific investigations involving laboratory testing and collecting information from other sources.
A few suggested projects are as follows:
Checking the bacterial contamination in drinking water by testing sulphide ion
Study of the methods of purification of water.
Testing the hardness, presence of Iron, Fluoride, Chloride, etc., depending upon the regional variation in drinking water and study of causes of presence of these ions above permissible limit (if any).
Investigation of the foaming capacity of different washing soaps and the effect of addition of Sodium carbonate on it.
Study the acidity of different samples of tea leaves.
Determination of the rate of evaporation of different liquids Study the effect of acids and bases on the tensile strength of fibres.
Study of acidity of fruit and vegetable juices.
Note: Any other investigatory project, which involves about 10 periods of work, can be chosen with theapproval of the teacher.
Practical Examination for Visually Impaired Students of Class 11
Below is a list of practicals for the visually impaired students.
A. List of apparatus for identification for assessment in practicals (All experiments) Beaker, tripod stand, wire gauze, glass rod, funnel, filter paper, Bunsen burner, test tube, test tube stand, dropper, test tube holder, ignition tube, china dish, tongs, standard flask, pipette, burette, conical flask, clamp stand, dropper, wash bottle • Odour detection in qualitative analysis • Procedure/Setup of the apparatus
B. List of Experiments A. Characterization and Purification of Chemical Substances 1. Crystallization of an impure sample of any one of the following: copper sulphate, benzoic acid B. Experiments based on pH 1. Determination of pH of some solutions obtained from fruit juices, solutions of known and varied concentrations of acids, bases and salts using pH paper 2. Comparing the pH of solutions of strong and weak acids of same concentration.
C. Chemical Equilibrium 1. Study the shift in equilibrium between ferric ions and thiocyanate ions by increasing/decreasing the concentration of eitherions. 2. Study the shift in equilibrium between [Co(H2O)6]2+ and chloride ions by changing the concentration of either of the ions.
D. Quantitative estimation 1. Preparation of standard solution of oxalic acid. 2. Determination of molarity of a given solution of sodium hydroxide by titrating it against standard solution of oxalic acid.
E. Qualitative Analysis 1. Determination of one anion and one cation in a given salt 2. Cations – NH+4 Anions – (CO3)2-, S2-, (SO3)2-, Cl-, CH3COO- (Note: insoluble salts excluded) 3. Detection of Nitrogen in the given organic compound. 4. Detection of Halogen in the given organic compound.
Note: The above practicals may be carried out in an experiential manner rather than recording observations.
We hope students must have found this information on CBSE Syllabus useful for their studying Chemistry. Learn Maths & Science in interactive and fun loving ways with ANAND CLASSES (A School Of Competitions) App/Tablet.
Frequently Asked Questions on CBSE Class 11 Chemistry Syllabus
Q1
How many units are in the CBSE Class 11 Chemistry Syllabus?
There are 9 units in the CBSE Class 11 Chemistry Syllabus. Students can access various study materials for the chapters mentioned in this article for free at ANAND CLASSES (A School Of Competitions).
Q2
What is the total marks for practicals examination as per the CBSE Class 11 Chemistry Syllabus?
The total marks for the practicals as per the CBSE Class 11 Chemistry Syllabus is 30. It includes volumetric analysis, content-based experiment, salt analysis, class record, project work and viva.
Q3
Which chapter carries more weightage as per the CBSE Syllabus for Class 11 Chemistry?
The organic chemistry chapter carries more weightage as per the CBSE Syllabus for Class 11 Chemistry.
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