Mole concept is the method used to express the amount of substance. This has been experimentally proving that one gram atom of any element, as well as one gram molecule of any substance, contains the same amount of entities.

The experimentally decided number is found to be 6.022137 × 10²³. After the discovery of the mole concept, the problem of finding absolute atomic masses of atoms was solved. It was so because the mole concept helps to count the number of atoms or molecules in a definite amount of the given substance. Let’s learn more about mole concept formulas and examples.

What is a Mole Concept?

Mole concept is known as the method used to express the amount of substance. A mole is defined as the amount of substance containing the same number of different entities (such as atoms, ions, and molecules) as the number of atoms in a sample of pure ¹²C weighing precisely 12 g.

Even a gram of any pure element contains a high amount of atoms. The mole connects a simple macroscopic feature (bulk mass) to a genuinely significant fundamental trait (number of atoms, molecules, etc.). One mole is also defined as the amount of a substance that contains as many entities as there are atoms in exactly 12 g of the ¹²C isotope.

It was found that the mass of one atom of carbon-12 element is equal to 1.992648 × 10^-23 g as measured by the mass spectrometer. Since one mole of Carbon-12 atom weighs 12 g, therefore, the number of atoms in it equals 12 g mol^-1 / 1.992648 × 10²³ g atom^-1 = 6.0221367 × 10²³ atoms mol^-1. The following formula may be used to calculate the number of moles of a chemical in a given pure sample:

n = N/N_A

Where n represents the number of moles of the chemical, N denotes the average number of fundamental units in the sample, and N_A represents the Avogadro constant.

Avogadro’s Number (N_A)

Avogadro’s number is 6.0221367 × 10²³. It is named Avogadro’s constant or number in honor of Amedeo Avogadro, a great pioneer in this field. It is denoted by N_A.

We can also say that 1 mole is the collection of 6.0221367 × 10²³ units of a substance. Here, the substance is atoms, molecules, or ions. The number of units that make up a mole has been established empirically to be 6.0221367 × 10²³.

This is known as the fundamental constant, also known as Avogadro’s number (N_A) or the Avogadro constant. This constant is appropriately stated in chemistry using an explicit unit termed per mole.

No matter what the given substance may be, one mole of it is always equal to N_A.

Mole Concept Formulas

Mole concept is a method where we identify the mass of chemical substances as per requirement.

The entire mole concept revolves around 12 g (0.012 kg) of the ¹²C isotope. In the SI system, the unit of the fundamental quantity ‘the amount of substance’ is the mole.

The symbol of the mole is “mol”. Below are the quantities related to the mole concept:

Atomic Mass and Molecular Mass

Atomic mass of an element is the mass of its one atom. The unit of atomic mass is a.m.u. The atomic mass is roughly the sum of protons and neutrons present in the atom. One atomic mass unit (a.m.u.) is said to be exactly equal to one-twelfth the mass of one carbon-12 atom. Therefore, the value of one a.m.u. is 1 g / N_A = 1.66056 × 10^-24 g. In the present era, the atomic mass unit is known as a unified mass unit. Hence, a.m.u. has been replaced by u. 

For example, the atomic mass of carbon is 12.011 amu since carbon mostly contains carbon-12 isotope. Carbon-13 is 1.1%, and very few traces of carbon-13. The atomic mass of these isotopes is different.

Molecular Mass is the sum of the masses of the atoms present in a given molecule. The unit of molecular mass is a.m.u. However, if the molecular mass of a mole of a substance is asked, then the unit used is grams despite the fact that the SI unit is kilograms. The molecular mass of a molecule is defined as the relative mass of its molecule when compared to the mass of a ¹²C atom divided into 12 units. In layman’s words, it denotes the number of times a molecule of the relevant material is heavier than an atom. 

For example, the molecular mass of water (H₂O) is 18.015 amu, the atomic mass of a hydrogen atom is 1.007 amu and the atomic mass of oxygen is 15.99 amu.

What is Molar Mass?

Molar mass is defined as the total mass of one mole of the substance. It is frequently expressed in terms of ‘grams per mole’ (g/mol). The SI unit for this amount, however, is kg/mol. The following formula may be used to calculate molar mass:

Molar mass of a Substance = (Mass of the substance in grams)/(Number of Moles)

The molar mass of water, for example, is roughly 18.015 g/mol, which is the mass of N_A number of water molecules.

Gram Atomic Mass and Gram Molecular Mass

Gram atomic mass is the mass of one mole of an atom. Gram molecular mass is the mass of one mole of a molecular substance expressed in grams. It is also known as molar mass. It is also defined as the mass of one mole of molecules. This is the difference between gram atomic and gram molecular mass. This amount of a substance is also called one gram molecule. 

How to Find Gram Molecular Mass?

A substance’s gram molecular mass is its molecular mass measured in grams. e.g., the molecular mass of O₂ is 32 grams; this is the relative molecular mass given in grams. Remember that relative atomic mass is always expressed as a ratio and has no units. 

Gram Molecular Volume

The volume occupied by one gram mole of a substance in a vapor state or gaseous state at STP is called as gram molecular volume. The standard temperature to obtain GMV is 273K, and the standard pressure is 1 atm. The ideal gas equation is utilized to get gram molar volume (GMV = 22.4L).

Relative Molecular Mass (RMM)

Relative molecular mass is the molecular weight of an element or molecule; it is expressed as RMM. It is the number of times a single molecule of a substance remains heavier than one-twelfth the mass of a carbon atom (¹²C).

Related Formulas

Following are the formulas for the number of moles, number of atoms and molecules, and the relationship between 1 amu and grams.

• Number of Moles = Mass of the Sample/Molar mass
• Number of atoms or molecules = Number of moles × (6.023 × 10²³)
• 1 amu = 1 gram/(6.023 × 10²³) = 1.66 × 10^-24 grams

Number of Electrons in a Mole of Hydrogen Molecule

One mole of H₂ contains 6.023 × 10²³ molecules, and each molecule of hydrogen contains two electrons.

1 mole = 6.023 × 10²³

Hence, the total number of electrons in 1 mole of H₂ is 12.046 × 10²³.

Solved Examples on Mole Concept

Example 1: Calculate the molecular mass of Ammonium Sulphate (NH₄)₂SO₄. 

Solution:

Since, the relative atomic masses of N = 14, H = 1, S = 32, O =16

Therefore, the molecular mass of the given compound is,

= 2 (14 × 1 + 1 × 4) + 32 + 16 × 4

= 2(14 + 4) + 32 + 64

= 2 × 18 + 32 + 64

= 36 + 32 + 64

= 132 amu

Example 2: The molecular mass of H₂SO₄ is 98 a.m.u. What does it mean?

Solution:

It means that one molecule of H₂SO₄ is 98 times as heavy as 1/12 the mass of a C-12 atom.

Example 3: Calculate the number of atoms present in 18g of H₂O.

Solution:

The 18 g of H₂O ⇒ 1 mole of H₂O ⇒ N_A molecules

Now, 1 molecule of H₂O contain 3 atoms

⇒ 1 mole H₂O will contain 3*N_A atoms = 3 × 6.022 × 10²³ atoms = 1.8066 × 10²⁴ atoms

Example 4: How many moles are present in 200g of NaOH?

Solution:

The mass of 1 mole of NaOH = 23 + 16 + 1= 40 g

Therefore, in 200g of NaOH the number of moles present = 200g / 40g mol^-1 = 5 mol

Example 5: Calculate the mass of an atom of oxygen element.

Solution:

Mass of 1 mole of oxygen = 16g

No. of atoms in 1 mole of oxygen = N_A

Therefore, mass of one atom of oxygen = 16g / N_A = 16 / (6.022 × 10²³) = 2.657 × 10^-23 g

Example 6: Find the ratio of moles of oxygen atoms present in the compounds H₂SO₄, H₂SO₃, and SO_2.

Solution:

1 mole of H₂SO₄ contains 4 × N_A atoms of oxygen

1 mole of H₂SO₃ contains 3 × N_A atoms of oxygen

1 mole SO₂ contains 2 × N_A atoms of oxygen 

Therefore the required ratio is 4 × N_A : 3 × N_A : 2 × N_A = 4 : 3 : 2

Example 7: How many moles of hydrogen and oxygen gas are required to produce 13 moles of water?

Solution:

Chemical equation of water formation:

H₂ + O₂ ⇢ H₂O

Now the balanced equation is

2H₂ + O₂ ⇢ 2H₂O

Thus we can deduce that 2 moles of hydrogen gas and 1 mole of oxygen gas combine together to form 2 moles of water. We can write the equation as :

H₂ (1 mole) + O₂ (1/2 mole)⇢ H₂O (1 mole)        

Hence, for production of 13 moles of water the required chemical equation would be:

H₂ (13 moles) + O₂ (13/2 moles) ⇢ H₂O (13 moles)

Hence, for production of 13 moles of water we need 13 moles of hydrogen and 6.5 moles of oxygen gas.

Example 8: Calculate moles of electrons present in 104 g of acetylene gas.

Solution:

Now, formula for acetylene gas is C₂H₂. It’s structure is H—C≡C—H. Therefore, number of electrons present in 1 molecule of acetylene are 14.

Now, mass of 1 mole of acetylene is 26g ⇒ 104g of acetylene are 4 moles.

Now, 1 mole of acetylene = N_A molecules

⇒ N_A molecules have 14 × N_A electrons

⇒ 4 moles of acetylene have 4 × 14 × N_A electrons = 56 × N_A electrons = 56 moles of electrons.

FAQs on Mole Concept

Question 1: What is the Mole concept?

Answer:

Mole concept is a method where we identify the mass of chemical substances as per requirement. The entire mole concept revolves around 12 g (0.012 kg) of the ¹²C isotope. In the SI system, the unit of the fundamental quantity ‘amount of substance’ is the mole. 

Question 2: What is one mole equal to?

Answer:

One mole is the collection of 6.0221367 × 10²³ units of a substance. Here, the substance is atoms, molecules, or ions. 

Question 3: How do we use mole fraction?

Answer:

Mole fraction is defined as the number of single component molecules divided by the total number of molecules. When two reactive elements are combined together, the ratio of the two elements is understood by the mole fraction.

Question 4: What is the importance of the mole concept?

Answer:

Mole concept provides a standard for measuring certain masses that are universally accepted. It provides a link between several masses and is an efficient way to measure macroscopic properties. For the study of chemistry, understanding mole concept is essential, and knowledge of how mole is applied to masses, number of substances, etc., is important.

Question 5: Is mole fraction equal to partial fraction?

Answer:

It can be said that the partial fraction of each gas is proportional to the mole fraction in a mixture.

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.

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CBSE Class 11 Chemistry Syllabus

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.

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‐ , S^2‐, NO₂^‐ , SO₃^2‐, SO^2‐ , NO ^‐ , Cl^‐ , Br^‐, I‐, PO₄^3‐ , 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 the approval 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.