In our daily lives, we witness several reactions such as iron rusting, paper burning, curd sourness, ozone generation, and so on. Many of these reactions require the presence of components in distinct phases, such as solid iron reacting with gaseous oxygen to generate solid iron oxide, which we call rust.
Similarly, gaseous hydrogen and gaseous oxygen combine to make liquid water. Having to deal with such reactions is a time-consuming chore. When the components are in the same phase, their interaction is simple to understand; but, when the components are in different stages, the interaction becomes more complicated.
Table of Contents
Homogeneous and Heterogeneous equilibrium
To simplify difficulties and grasp the notion, we divide such reactions into two categories: homogeneous reactions, in which the components involved are present in the same phase, and heterogeneous reactions, in which the components involved are present in separate phases. The methods for dealing with both reactions, as well as the determination of the equilibrium state, differ.
The reaction in which all of the products and reactants have the same phase. For instance, all of the products and reactants could be gases or all of the products and reactants could be liquids. An equilibrium reaction is one that can be reversed and forwarded while the concentrations of the reactants and products remain constant.
Equilibrium is a chemical reaction state in which the rate of forwarding and backward reaction is the same. Furthermore, there are two forms of equilibrium: homogeneous equilibrium and heterogeneous equilibrium.
A homogeneous equilibrium in one phase is defined as a homogeneous mixture (reactants and products in a single solution). Remember that the reactants are on the left side of the equation and the products are on the right. As a result, the reaction between the solutes corresponds to a single homogeneous equilibrium.
A heterogeneous equilibrium, on the other hand, is a reaction system in which the products and reactants exist in two or more phases.
A homogeneous equilibrium can be further classified into two types. The number of molecules in the product in the first category is the same as the number of molecules in the reactants in that particular equation. As an example:
N2 (g)+O2 (g) ⇌ 2NO (g)
From the above example, we can see that there are two molecules of reactants (one of each) and two molecules of product on the right side. In the second category of a homogeneous equilibrium equation, the opposite events occur. The product’s molecule count is not the same as or equal to the reactant’s molecule count. As an example:
2SO2 (g)+O2 (g) ⇌ 2SO3(g)
We can see from the preceding example that there are only three molecules of reactant and two molecules of product present in the reaction. The reactions in liquid solutions between solutes belong to one type of homogeneous equilibria in a homogeneous equilibrium, and the chemical species involved can be molecules, ions, or a mixture of both.
Difference between KC and KP
The two equilibrium constants are distinguished by the fact that they are applied to different concentrations. KP denotes the equilibrium constant at partial pressure during a reaction. These values can be calculated using the reactant and product values, the equations, and the specific values of those formulae. A relationship exists between the two equilibrium constants, which is shown below:
KP = KC(RT)Δn
Difference Between Homogeneous and Heterogeneous Equilibrium
Equilibrium is defined as a situation in which the concentrations of reactants and products are constant. Homogeneous equilibrium and heterogeneous equilibrium are the two types of equilibria. The primary distinction between homogeneous and heterogeneous equilibrium is that in homogeneous equilibrium, the reactants and products are in the same phase of matter, but in heterogeneous equilibrium, they are in different phases.
Furthermore, when determining the equilibrium constant for homogeneous equilibria, we must include the concentrations of all reactants and products; however, when determining the equilibrium constant for heterogeneous equilibria, we must exclude the concentrations of solids and pure liquids and use the concentrations of other reactants and products. As an illustration,
2SO2(g) + O2(g) ⇌ 2SO3(g) is a homogeneous equilibrium, and
O2(g) + 2C(s) ⇌ 2CO(g) is an example for a heterogeneous equilibrium.
To put it simply, equilibrium is a situation in which the concentrations of reactants and products remain constant. Homogeneous equilibrium and heterogeneous equilibrium are the two types of equilibria. The primary distinction between homogeneous and heterogeneous equilibrium is that in homogeneous equilibrium, the reactants and products are in the same phase of matter, but in heterogeneous equilibrium, they are in different phases. Furthermore, the equilibrium constant for homogeneous equilibria includes all reactant and product concentrations, whereas the equilibrium constant for heterogeneous equilibria must omit solid and pure liquid concentrations.
FAQs
Question 1: What is meant by heterogeneous equilibria?
Answer:
The heterogeneous equilibrium refers to a system in which the reactants and products exist in two or more phases. The system’s phases are any combination of liquids, gases, solids, and solutions. It is vital to remember that pure liquids and solids cannot appear as equilibrium constant expressions when dealing with heterogeneous equilibrium.
Question 2: What is meant by buffer solutions?
Answer:
Buffer solutions are composed of either a weak base and its conjugate acid or a weak acid and its conjugate base. When extra ions are introduced to a buffer solution, the pH of the solution changes. According to Le Chatelier’s principle, when more ions are introduced, the equilibrium changes and the reactions shift to favour the solid or deionized form. In the case of an acidic buffer, the concentration of the hydrogen ion reduces, and the solution produced is less acidic than a solution containing pure weak acid.
Question 3: What is an equilibrium constant?
Answer:
The word equilibrium constant can be defined as the expression that reflects the concentration of reactants and products after the chemical process has reached equilibrium. Temperature is critical in maintaining the equilibrium constant within the reactions. If the temperature remains constant, the equilibrium remains constant as well. This is evident throughout the equation, and it ultimately plays a critical role in maintaining a constant balance.
Question 4: What is the difference between a homogeneous mixture and a heterogeneous mixture?
Answer:
When viewed at a macroscopic level, homogeneous mixes are frequently thought to be indistinguishable from the pure substance. The reaction that occurs between the solutes is part of a single homogeneous equilibrium. Sugar, salt, water, dye, air, and blood are examples of homogeneous mixes. A heterogeneous mixture has a distinct identifying quality in which the various components of the combination can be seen. It is a reaction system in which the reactants and products are found in two or more phases. Pizza, cookies, rocks, and other such items are examples of heterogeneous mixtures.
Question 5: What is meant by the common ion effect and its role?
Answer:
The common ion effect depicts the changes that occur when ions are injected into a solution containing the same ion. When common ions are added to a solution, the solubility of a molecule decreases due to a shift in the equilibrium. The common ion effect is crucial in the modulation of buffers. A buffering solution contains either an acid or a base, either of which is accompanied by its conjugate counterpart. The pH of the solution changes as additional conjugate ions are added. This effect must be considered when evaluating solution equilibrium when common ions are introduced.
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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