Grade 11
  1. Basic concepts of chemistry  1.1. Importance of Chemistry  1.2. Nature and scope of chemistry  1.3. laws of chemical combination  1.3.1. Law of conservation of mass  1.3.2. Law of definite proportions  1.3.3. Law of multiple proportions  1.3.4. Law of gaseous volume  1.3.5. Avogadro's Law  1.4. Dalton's atomic theory  1.5. Mole concept and molar mass  1.6. Percent composition  1.7. Empirical and molecular formula  1.8. Stoichiometry and Stoichiometric Calculations  1.9. Limiting reagent concept
  2. Structure of the atom  2.1. Discovery of the electron, proton, and neutron  2.2. Atomic Model  2.2.1. Thomson's model  2.2.2. Rutherford's model  2.2.3. Bohr's model  2.3. Quantum mechanical model of the atom  2.4. Dual nature of matter and radiation  2.5. Heisenberg's uncertainty principle  2.6. Quantum numbers  2.7. Atomic orbitals and their shapes  2.8. Electronic configuration of elements  2.9. Aufbau principle  2.10. Pauli's exclusion principle  2.11. Hund's maximum multiplicity rule  2.12. de Broglie's hypothesis  2.13. Photoelectric effect
  3. Classification of elements and periodicity in properties  3.1. Historical development of the periodic table  3.2. Modern Periodic Law and Periodic Table  3.3. Periodic trends in properties  3.3.1. Atomic and Ionic Radius  3.3.2. Ionization enthalpy  3.3.3. Electron gain enthalpy  3.3.4. Electronegativity  3.3.5. Valency  3.3.6. Metallic and Non-Metallic Properties  3.3.7. Oxidation states  3.4. Unusual Properties of the First Elements
  4. Chemical Bonding and Molecular Structure  4.1. Kossel–Lewis approach to chemical bonding  4.2. Ionic or electrovalent bond  4.3. Covalent bond and its features  4.4. Bond parameters  4.5. VSEPR Theory  4.6. Valence bond theory  4.7. Hybridization and its types  4.8. Molecular orbital theory  4.8.1. Bond order and stability  4.9. Hydrogen bonding
  5. States of matter  5.1. Intermolecular forces  5.2. Gas Laws  5.2.1. Boyle's law  5.2.2. Charles's Law  5.2.3. Avogadro's Law  5.2.4. Dalton's law of partial pressure  5.2.5. Graham's law of spread  5.3. Ideal Gas Equation and Applications  5.4. Real gases and deviations from ideal behaviour  5.5. Kinetic molecular theory of gases  5.6. Liquefaction of gases  5.7. Properties of Liquids  5.8. Surface tension and viscosity
  6. Thermodynamics  6.1. System and environment  6.2. Types of systems in thermodynamics  6.3. Types of procedures  6.4. First law of thermodynamics  6.5. Internal energy and enthalpy  6.6. Heat capacity and specific heat  6.7. Hess's law of constant heat summation  6.8. Bond dissociation enthalpy  6.9. Enthalpy of formation and combustion  6.10. Enthalpy of solution and neutralization  6.11. Spontaneity and the second law of thermodynamics  6.12. Gibbs free energy and its applications
  7. Balance  7.1. The concept of balance  7.2. Equilibrium constant and its applications  7.3. Le Chatelier's principle  7.4. Ionic equilibrium in solutions  7.5. Acid and Base Theory  7.5.1. Arrhenius concept  7.5.2. Bronsted-Lowry concept  7.5.3. Lewis concept  7.6. pH Scale and pOH  7.7. Common ion effect  7.8. Hydrolysis of salts  7.9. Buffer solutions and their mechanism of action  7.10. Solubility equilibrium of sparingly soluble salts
  8. Redox reactions  8.1. Oxidation and reduction concepts  8.2. Oxidation number and its applications  8.3. Redox reactions in terms of electron transfer  8.4. Balancing redox reactions (oxidation number and ion-electron methods)  8.5. Types of redox reactions  8.6. Electrochemical Cells and Redox Reactions
  9. Hydrogen  9.1. Position of Hydrogen in the Periodic Table  9.2. Isotopes of hydrogen  9.3. Preparation of Hydrogen  9.4. Properties of Hydrogen  9.5. Hydrides and their classification  9.6. Water and heavy water  9.7. Hydrogen peroxide and its properties  9.8. Uses of Hydrogen and Its Compounds
  10. S-block elements (alkali and alkaline earth metals)  10.1. Group 1 Elements - Properties and Trends  10.2. Group 2 Elements - Properties and Trends  10.3. Unusual properties of lithium and beryllium  10.4. Important compounds of sodium and their uses  10.5. Important Compounds of Magnesium and Calcium
  11. Some p-block elements  11.1. General Introduction to p-Block Elements  11.2. Group 13 Elements  11.2.1. Boron and its compounds  11.3. Group 14 Elements  11.3.1. Carbon and its allotropes  11.3.2. Important Compounds of Carbon and Silicon
  12. Organic Chemistry - Some Basic Principles and Techniques  12.1. Classification and Nomenclature of Organic Compounds  12.2. Structural representation of organic compounds  12.3. Classification of organic reactions  12.4. Electronic Effects in Organic Chemistry  12.4.1. Inductive effect  12.4.2. Resonance effect  12.4.3. Hyperconjugation  12.5. Reaction mechanism and intermediate species  12.6. Purification and qualitative analysis of organic compounds
  13. Hydrocarbons  13.1. Hydrocarbons  13.1.1. Preparation and Properties of Alkenes  13.2. alkene  13.2.1. Preparation and Properties of Alkenes  13.2.2. Electrophilic addition reactions  13.3. Alkynes  13.3.1. Preparation and Properties of Alkynes  13.4. Aromatic hydrocarbons  13.4.1. Structure and properties of benzene  13.4.2. Electrophilic substitution reactions of benzene
  14. Environmental Chemistry  14.1. Environmental Pollution  14.2. Atmospheric pollution and the greenhouse effect  14.3. Water pollution and treatment methods  14.4. Soil pollution and its effects  14.5. Industrial waste and its treatment  14.6. Strategies for pollution control

Grade 11 → Basic concepts of chemistry ↓


Mole concept and molar mass


The mole concept is a fundamental principle in chemistry that serves as a bridge between the atomic world and the macroscopic world that we can measure. Understanding this concept is important for further studies in chemistry and is widely used in various calculations and reactions.

What is a mole?

The mole is a unit of measurement used in chemistry to express the amount of a chemical substance. It is one of the seven base SI units and is defined as exactly 6.02214076 × 10 23 particles, which may be atoms, molecules, ions, or electrons. This number is known as Avogadro's number.

Importance of mole

  • Counting particles: Atoms and molecules are incredibly small. They are so small that it would be impractical to count them one by one. The mole allows chemists to count particles by weighing them.
  • Chemical reactions: Chemical reactions often involve very large numbers of atoms and molecules, and a convenient unit is needed to deal with them. The mole allows chemists to measure substances in quantities that correspond to the amount needed for a chemical reaction.
  • Universal standard: Just as a dozen is universally accepted as 12, a mole is universally accepted as 6.02214076 × 10 23 particles.

Avogadro number: Visual representation

Let's understand Avogadro's number with an example for better understanding. Imagine you have a mole of basketballs.

One mole of basketballs (6.022 × 10 23 basketballs)
1 sesame seed

The number of basketballs you would have would be more than enough to cover the entire Earth with basketballs, piling them up to over a mile high. This helps us understand just how large Avogadro's number is.

Converting moles to particles and vice versa

An important skill in chemistry is to convert between moles and numbers of particles (atoms, molecules, etc.). To do this, we use Avogadro's number.

Conversion formulas

  • Number of particles = moles × Avogadro's number
  • Mole = Number of particles ÷ Avogadro's number

Example calculation

Suppose you have 2 moles of water molecules. How many water molecules do you have?

Number of water molecules = 2 moles × 6.022 × 10 23 molecules/mole = 1.2044 × 10 24 molecules

Now, let's do the calculation in reverse. How many moles are there 1.2044 × 10 24 water molecules?

Moles of water = 1.2044 × 10 24 molecules ÷ 6.022 × 10 23 molecules/mole = 2 moles

Molar mass

Molar mass is the mass of one mole of a substance (element or compound) expressed in grams per mole (g/mol). It is numerically equal to the average atomic mass of the element in atomic mass units (amu).

How to calculate molar mass

To find the molar mass of a molecule, add up the atomic masses of all the atoms that make up that molecule.

Example: Molar mass of water (H 2 O)

A water molecule consists of two hydrogen atoms and one oxygen atom. To calculate the molar mass, follow these steps:

  1. Mass of hydrogen (H) = 1.01 g/mol
  2. Mass of oxygen (O) = 16.00 g/mol
  3. Molecular Formula: H 2 O = 2(H) + 1(O)
Molar Mass of H 2 O = 2(1.01) + 16.00 = 2.02 + 16.00 = 18.02 g/mol

Filling the concept with examples

If you have 36.04 grams of water, how many moles will you have?

Moles of water = Mass of water ÷ Molar Mass of water = 36.04 grams ÷ 18.02 g/mol = 2 moles

This means you have 2 moles of water molecules.

Mole-mass-number relation

Chemistry often requires you to convert between number of particles, moles, and mass. Here's how these concepts are related:

Conversion flow

Let's say you want to convert mass to moles or particles. Think of it as a flowchart:

Mass (grams) Moles Number of particles

To convert between these, you can use the following:

  • Mass to moles: Divide the mass by the molar mass
  • Moles to particles: Multiply the number of moles by Avogadro's number
  • Particles to moles: Divide the number of particles by Avogadro's number
  • Moles to mass: Multiply the number of moles by the molar mass

Custom question

Question 1

If you have 4 moles of glucose (C 6 H 12 O 6), what is the mass in grams?

Atomic masses: C = 12.01, H = 1.008, O = 16.00 Molar Mass of C 6 H 12 O 6 = 6(12.01) + 12(1.008) + 6(16.00) = 72.06 + 12.096 + 96.00 = 180.16 g/mol Mass = Moles × Molar Mass = 4 × 180.16 = 720.64 grams

Question 2

How many molecules are there in 0.5 moles of nitrogen gas (N2)?

Number of molecules = Moles × Avogadro's number = 0.5 × 6.022 × 10 23 = 3.011 × 10 23 molecules

These examples help in understanding the mole concept and its usefulness in chemistry.

Conclusion

The mole concept and molar mass are central to understanding and performing calculations in chemistry. They provide a way to convert atomic-scale interactions into macroscopic quantities that we can measure and manipulate in a laboratory setting. Mastering these concepts is essential for anyone who wishes to pursue further studies or a career involving chemistry.


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Mole concept and molar mass

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