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 ↓


laws of chemical combination


The laws of chemical combination are fundamental rules that explain how elements and compounds combine to form new substances. These rules were formulated based on systematic experiments and are the cornerstone of modern chemistry. There are several major laws associated with chemical combinations, and each law gives information about how substances interact at the molecular level.

Law of conservation of mass

The law of conservation of mass, established by Antoine Lavoisier in the late 18th century, states that mass in an isolated system is neither created nor destroyed by chemical reactions. In simple terms, the mass of the reactants in a chemical reaction is equal to the mass of the products. This principle is fundamental to understanding chemical processes.

Consider the reaction of hydrogen gas H 2 and oxygen gas O 2 to form water H 2 O:

2H 2 + O 2 → 2H 2 O

If we start with 4 grams of hydrogen and 32 grams of oxygen, the total mass will be 36 grams. After the reaction, 36 grams of water are formed, respecting the law of conservation of mass.

Law of definite proportions

The law of definite proportions, also called the law of constant composition, was proposed by Joseph Proust in 1799. This law states that the constituent elements in a chemical compound are always present in a fixed proportion by mass, regardless of the source or method of formation.

For water, the ratio of the mass of hydrogen to the mass of oxygen is always 1:8. If you have a sample with 2 grams of hydrogen, it will combine with 16 grams of oxygen to form 18 grams of water.

Mass ratio in water = 1 part hydrogen : 8 parts oxygen

Law of multiple proportions

The law of multiple proportions was formulated by John Dalton in 1804. This law states that if two elements combine to form more than one compound, then the ratio of the weight of the element that combines with the fixed weight of the other element are simple whole numbers.

An example of this is found in the compounds carbon monoxide CO and carbon dioxide CO 2. Here, 12 grams of carbon can combine with 16 grams of oxygen to form carbon monoxide and 32 grams of oxygen can combine with 32 grams of oxygen to form carbon dioxide. The ratio of the mass of oxygen combined with a constant mass of carbon (12 grams) is 1:2.

Mass ratio of oxygen in CO : CO 2 = 16 : 32 = 1 : 2

Law of reciprocal proportions

The law of reciprocal proportions, sometimes called the law of equivalent proportions, was proposed by Jeremias Richter in 1792. This law states that the masses of elements that separately combine with a fixed mass of another element are the same as the masses of the elements that combine directly with each other or with simpler multiples of it.

Consider the elements hydrogen H, oxygen O, and sulfur S:

Hydrogen combines with oxygen to form water, and with sulfur to form hydrogen sulfide. If water and hydrogen sulfide are formed via combination steps by stabilizing the respective masses of oxygen and sulfur against the same mass of hydrogen, the ratio is consistent.

Mass ratio of oxygen in H 2 O : Mass ratio of sulphur in H 2 S = 8:16

Gay-Lussac's law of gaseous volumes

In 1808, Joseph Louis Gay-Lussac proposed that when gases react together at a constant temperature and pressure, the volumes of the gaseous reactants and products are in the ratio of small whole numbers. This law is particularly useful when dealing with gaseous substances.

For example, ammonia is made by the reaction of nitrogen and hydrogen gases:

N 2 + 3H 2 → 2NH 3

One volume of nitrogen reacts with three volumes of hydrogen to form two volumes of ammonia, which represents a simple whole number ratio.

Practical applications

These rules are applied in a variety of chemical calculations and reactions:

  • Stoichiometry: The quantitative relationship between reactants and products in a chemical reaction is derived from these rules.
  • Preparing chemical equations: Chemical equations are balanced using the law of conservation of mass to ensure that matter is neither created nor destroyed.
  • Formula determination: The law of definite proportions is used to determine the empirical and molecular formulas of compounds.
  • Understanding compounds: Rules such as multiple proportions help distinguish between compounds with the same elements.
  • Gaseous reactions: Using Gay-Lussac's law, chemists can predict the outcomes of reactions involving gases.

Historical influences

The laws of chemical combination were important in the development of chemistry as a scientific discipline. They helped establish the atomic theory, which holds that matter is composed of atoms. Together with experimental evidence, these laws provided a systematic approach to understanding chemical reactions, thus laying the foundations of modern chemistry.

Before these rules, chemistry, which was mainly based on mysticism and pseudoscience, dominated the understanding of chemical processes. By incorporating quantitative methods and these rules, chemistry was transformed into a rigorous and systematic science. These rules paved the way for researchers to predict and quantify reactions with greater accuracy, leading to discoveries and advances in various chemical industries and scientific investigations.


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