Jeremias Richter, a German chemist, was the first to create or discover the word Stoichiometry.
The quantitative analysis of the reactants and products involved in a chemical reaction is known as chemical stoichiometry.
The name “stoichiometry” comes from the Greek words “stoikhein” (element) and “metron” (measure).
We will study what it implies and discuss the various components of this topic further below.
Table of Contents
What Is Stoichiometry?
The calculation of products and reactants in a chemical reaction is known as stoichiometry. It is a key concept in chemistry that allows us to compute reactant and product amounts using balanced chemical equations. The balanced equation’s ratios are used in this case. In general, all reactions are influenced by one fundamental factor which is the amount of substance present.
Stoichiometry is the measurement of the products and reactants in any chemical reaction. The term “stoichiometry” is derived from two Greek words: “stoichion,” which refers to element determination, and “metry,” which refers to measurement.
Moreover, stoichiometry is founded on the law of conservation of mass, which states that the total mass of reactants equals the total mass of products, proving that product and reactant numbers are frequently stated as a positive integer ratio. This illustrates that if you know the number of specific products, calculating the product amount is simple. The computation of other reactants is also possible if the quantity of one reactant is known and the resultant quantity can be calculated via an experiment.
Stoichiometry is a measurement of the quantitative relationship between the products and reactants of a chemical process expressed in mass or volume ratios. Stoichiometry is a fundamental mathematical principle that describes the rule of conservation of mass, which states that matter cannot be generated or destroyed, but can only be transformed from one condition to another. A balanced stoichiometry is required for a chemical reaction to occur and progress to completion. In a chemical reaction, the quantity of each chemical element on the product side should be equal to the corresponding element quantity on the reactant side. Stoichiometry assists in determining the amount of substance required or present. The following are characteristics that can be measured:
The mass of reactants and products
Molecular mass
A chemical equation, states the reactant side.
Chemical equations
What is the Stoichiometric Coefficient?
The number of molecules involved in a process is known as the stoichiometric coefficient or stoichiometric number. When you look at a balanced reaction, you’ll see that both sides of the equation have the same amount of elements. The number in front of atoms, molecules, or ions is known as the stoichiometric coefficient. Fractions and whole numbers can both be used as stoichiometric coefficients. The coefficients essentially assist us in determining the mole ratio between reactants and products.
Both sides of the equation have the same amount of components in a balanced reaction. The stoichiometric coefficient is the number written in front of atoms, ions, and molecules in a chemical reaction to balance the number of each element on both the reactant and product sides of the equation. While fractions can be used as stoichiometric coefficients, whole numbers are more commonly utilized and preferred. Since they establish the mole ratio between reactants and products, these stoichiometric coefficients are useful. The Stoichiometric coefficient is the number of molecules of a reactant that participate in a reaction.
Consider the following equation:
aA + bB ⇌ cC + dD
The Stoichiometric coefficients of the A, B, C, and D, respectively, are a, b, c, and d in this equation.
Stoichiometry in Chemical Analysis
Chemists frequently utilize stoichiometric calculations, which follow a quantitative analysis approach, to quantify the amounts of chemicals present in a sample. There are two primary forms of analysis. We’ll go over them in more detail later.
Gravimetric Calculation-In analytical chemistry, gravimetric analysis refers to the quantitative determination of analyte based on the mass of the solid. The gravimetric analysis produces the most precise results of any other analytical method since a substance’s weight can be measured with more precision than other fundamental quantities.
The following are several types of gravimetric analysis.
Precipitation gravimetry – It entails isolating an ion in solution by a precipitation reaction, filtering, washing the precipitate to remove impurities, and finally weighing and measuring the precipitate’s mass by difference.
Volatilization gravimetry – The process of separating components of a mixture by heating or chemically decomposing the sample is known as volatilization gravimetry.
Electrogravimetry – It comprises the electrochemical reduction of metal ions at the cathode as well as the deposition of ions on the cathode at the same time. The mass of analyte initially present in the sample is determined by weighing the cathode before and after electrolysis, and the weight difference corresponds to the mass of analyte initially present in the sample.
Volumetric Analysis- Volumetric analysis is the process of quantifying a substance in terms of volume. In volumetric analysis, a known volume (V1) of a material whose concentration (N1) is known reacts with an unknown volume (V2) of a solution of a substance whose concentration (N2) must be predicted. At the end of the response, the volume, V1, is recorded. The following equation is used to compute the N2 concentration.
N2 x V2 = N1 x V1
Volumetric analysis terms include:
Titration – Titration is the process of determining the volume of solution required to thoroughly react with the volume of another solution.
Indicator – Indicators are chemicals that change color as the reaction progresses.
Titrant – A titrant is a solution with a known strength.
Titrate – The solution whose concentration is to be estimated is referred to as titrate.
Stoichiometric Calculations
It’s essential to solve stoichiometric problems. It necessitates an understanding and use of the mole idea, as well as the balancing of chemical equations and the careful conversion of units. Chemical equation-based difficulties can be categorized as.
Mole to mole relationships- In these problems, the moles of one of the reactants or products are computed if the moles of the other reactants and products are given.
Mass-Volume relationship- The mass or volume of one of the reactants or products is estimated using the mass or volume of other substances in these problems.
Mass-mass relationship- The mass of one of the reactants or products must be computed if the mass of the other reactants or products is given in these problems.
Volume-volume relationship- The volume of one of the reactants or products is given in these problems, while the volume of the other must be estimated.
Sample Questions
Question 1: Calculate how much sodium hydroxide will be needed to make 500mL of a 0.10 M solution.
Answer:
The molar mass of NaOH = 40g
Volume of NaOH= 500ml = 0.5 L
Molarity = 0.10M
Molarity = moles / volume in litres
So, weight of NaOH = molar mass of NaOH x volume x molarity
= 0.10 x 40 x 0.5
= 2 g
Question 2: What are Stoichiometric Calculations, and how do they work?
Answer:
Stoichiometric calculations entail a number of steps, including solving a balancing equation, determining the moles of chemicals produced in a reaction, and converting units to moles and vice versa using various conversion factors.
Question 3: To make 3 M 400 ml HCl, how much 11M HCl should be diluted with water?
Answer:
Given that, M1 = 11M, and M2 = 3M
Also, V1 = ? , anx V2= 400ml
Now, M1 x V1 = M2 x V2
V1 = (3×400) / 11
= 109 ml
Question 4: What is the Application of Stoichiometry in Real Life?
Answer:
Stoichiometry is used in the manufacturing of soaps, gasoline, tires, deodorants, fertilizers, and other products.
Question 5: By reacting nitrogen with hydrogen, how many moles of nitrogen are required to make 8.2 moles of ammonia?
Answer:
The balanced chemical equation is,
N2 + 3H2 → 2NH3
2 mole of NH3 are produced from = 1 mole of N2
8.2 mole of NH3 are produced from = (1/2) x 8.2
= 4.1 mol of N2
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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