Laws of Chemical Combination are one of the most fundamental building blocks of the subject of chemistry.
As in our surrounding different matter reacts with each other and form various kind of different substances.
Laws of Chemical Combination are the collection of laws that explains how these substances combine together to form anything at all. When matter reacts with another matter, a chemical reaction happens which changes the form, properties, or characteristics of the matter drastically.
This interaction of matter with each other is governed by the Laws of Chemical Combination.
Table of Contents
What is Law of Chemical Combination
A collection of laws that governs the interaction; such as how they combined to form other matter, of matter with each other, is combinedly called the Laws of Chemical Combinations. The law of chemical combination collection includes five laws, which are as follows:
Laws of Conservation of Mass
Laws of Definite Proportions
Laws of Multiple Proportions
Gay-Lussac’s Law of Gaseous Volumes
Avogadro’s Law
Law of Conservation of Mass
Law of Conservation of Mass states that “Mass can neither be created nor can be destroyed in a chemical reaction” but it can be transformed from one form to another.
In other words, in an enclosed system whenever matter undergoes a chemical or physical change, the total mass of reacting matter (reactants or matter before the change) is exactly equal to the total mass of reaction products. As no products or reactants are allowed to exit the system in a closed system, there is no loss of the substance, and hence the mass is conserved.
For example, in automobiles, fuel is burned to create energy for vehicles to move forward and the fuel is converted into fumes such as carbon dioxide, carbon monoxide, or sulfur dioxide, etc. This law was first outlined by Mikhail Lomonosov in 1756 and then further refinements are proposed by Antoine Lavoisier in 1773 after a lot of experiments.
Example of Law of Conservation of Mass
For an example of conservation of mass, let’s consider the formation of a water molecule from oxygen and hydrogen molecules. Balance chemical Reaction of formation of water is:
2H2 + O2 → 2H2O
In the above example, 2 molecules of Hydrogen combine with 1 molecule of Oxygen to form 2 molecules of water.
Mass of H is 1 unit and O is 16 units.
Mass of 2 molecules of Hydrogen (2H2) is 4 units.
Mass of 1 molecule of Oxygen (O2) is 32 units.
Mass of 2 molecules of water (2H2O) is 36 units.
2H2 + O2 → 2H2O
4 unit 32 unit 36 unit
The total mass of the reactants equals the total mass of the products in this case. Furthermore, the number of hydrogen and oxygen atoms in the reactant and product sides are also equal.
Law of Definite Proportions
Law of Definite Proportions which is also known as the Law of Constant Composition, states that in any given chemical compound the composition of the element by mass is always remains the same.
In other words, this means that the ratio of the mass of the elements in the chemical compounds always remains the same i.e., in a water molecule there will always be two hydrogens and one oxygen molecule. The law of Definite proportion is first proposed by Joseph Louis Proust in the late 18th century and many scholars have proved it since then. This law helps us identify between different chemical compounds as well.
Example of Law of Definite Proportions
Consider the different molecules of Oxides of Nitrogen for examples of the law of definite proportions.
In the above example of oxides of Nitrogen if we take any amount of Oxygen and Nitrogen in the sample then the ratio of nitrogen to oxygen in the formed NO will always be 1:1. Similarly, the ratio of nitrogen to oxygen for NO2, N2O, and N2O2 will be 1:2, 2:1 and 2:2 respectively. This is guaranteed by the law of definite proportions.
Law of Multiple Proportions
According to the Law of Multiple Proportions, If two elements combined to form more than one compound under different circumstances, then the ratio of the masses for one element when second mass is fixed for all different compounds is always a small whole number.
The Law of Multiple Proportions is also known as Dalton’s Law, as it was first proposed by Dalton in the year 1804. This law doesn’t hold for non-stoichiometric compounds as well as heavy molecules such as polymers and oligomers.
Example of Law of Multiple Proportions
Carbon and oxygen combine to form two distinct compounds (under different circumstances). The first is the most common gas, CO2 (Carbon dioxide), and the second is CO (Carbon monoxide).
Lets take 12 grams of carbon, and by the calculation of moles we can find that it reacts with 16 grams of Oxygen to make Carbon Monoxide and with 32 grams of oxygen to form Carbon dioxide.
As a result, the ratio of mass of oxygen in the first and second compounds is 2:1 = 32/16 = 2, (whole number).
Gay Lussac’s Law of Gaseous Volumes
Gay Lussac enacted this law based on his observations in 1808. This law states that “when gases are produced or combined in a chemical reaction, they do so in a simple volume ratio provided that all the gases are at the same temperature and pressure.”
This law is regarded as an of definite proportions for gases and the difference between these two chemical combination laws is that Gay Lussac’s Law is stated the ratio of volume, whereas the law of definite proportions is stated in terms of mass.
Example of Gay Lussac’s Law of Gaseous Volumes
In the above example 2 volumes of H2 combines 1 volume of O2 to form 2 volumes of H2O.
H2 (g) + O2 (g) → 2H2O (g)
Avogadro’s Law
According to Avogadro’s Law, an ‘equal volume of all gases contains the equal number of molecules under the same conditions of temperature and pressure.’
This law was proposed In 1811 by none other than Avogadro himself. In other words, this law states that the volume and number of moles of any gas are always directly proportional to each other. This means that two liters of hydrogen have the same number of molecules as two liters of oxygen at the same temperature and pressure.
Example of Avogadro’s Law
Equivalent volumes of different gases contain the same number of molecules at the same temperature and pressure. In the above example CL2 and H₂ has 1 volume each combines to form 2 volume of HCL.
Here,
Mole is a unit of measurement for substance. 1 mole substance contains 6.02214076×10²³ particles.
FAQs on Laws of Chemical Combination
What are the Laws of Chemical Combination?
The collection of laws which explains how the elements react with each other to form compounds is called the law of chemical combination and this collection includes Laws of Conservation of Mass, Laws of Definite Proportions, Laws of Multiple Proportions, Gay-Lussac’s Law of Gaseous Volumes, and Avogadro’s Law.
What is the Law of Conservation of Mass?
Law of Conservation of Mass states that “mass can neither can created not be destroyed, it can be converted from one form to another.”
What is the Law of Definite Proportions?
According to the Law of Definite Proportions, mass of the constituents of any given compounds always remains the same.
What are the limitations of the Law of Definite Proportions?
The Law of Definite Proportions doesn’t hold true for non-stoichiometric compounds such as Iron Oxide, whose chemical formula is generally written FeO but actually, it is Fe0.95O.
What is the Law of Multiple Proportions?
According to the Law of Multiple Proportions when a combination of two or more elements make more then two compositions, then the ratio of the mass of the compound of one element when the other remains fixed for any two compositions is always a small whole number.
Why are the laws of chemical combination important?
These laws provide an understanding of how the atoms of different elements combine to form compounds with various properties. Thus, these laws are important for the quantitative analysis of chemical systems.
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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