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 ↓


Law of multiple proportions


The law of multiple proportions is an essential principle in chemistry that was first articulated by John Dalton in the early 19th century. This law is fundamental to understanding chemical combinations and how elements interact to form different compounds. In its simplest form, this law states that if two elements can combine in different ways to form more than one compound, then the ratio of the masses of the second element that combine with a given mass of the first element will be a ratio of small whole numbers.

Understanding the concept

The main idea behind the law of multiple proportions is to understand how elements combine to form different chemical compounds. For this law to apply, more than one compound must be formed by two given elements. The key here is the concept of definite mass ratios, which are always simple whole numbers.

Let us break it down into simpler terms for better clarity. Consider elements A and B that form several compounds. According to the rule, the different masses of element B that combine with a fixed mass of element A will always be in the ratio of small whole numbers. This suggests a distinct and quantitative pattern in the way the elements combine.

Example of the law of multiple proportions

To see this rule in action, let's consider the classic example involving carbon and oxygen. These elements can form two compounds, namely carbon monoxide (CO) and carbon dioxide (CO2). We will illustrate this with a visual diagram.

    2 C + O 2 → 2 CO
    C + O 2 → CO 2

Now, let's look at the mass ratio:

  • In CO, 12g of carbon combines with 16g of oxygen.
  • In CO2 12g of carbon combines with 32g of oxygen.

Now, consider the mass ratio of oxygen that combines with the same mass of carbon (12 grams):

The ratio of oxygen to CO2 in CO is 16:32, which in simpler form is 1:2.

12 g carbon 16 grams of oxygen 12 g carbon 32 grams of oxygen

Importance of the law

The law of multiple proportions underlies the atomic theory proposed by Dalton, which states that atoms are indivisible units that form compounds in definite proportions. It emphasizes that chemical compounds are composed only of whole atoms, which corresponds to how the elements combine in a precise and consistent manner, forming a solid foundation for the development of modern chemistry.

This law was a strong evidence for Dalton's atomic theory, which supported the idea that chemical reactions do not involve the transformation of substances into new types of atoms but rather the recombination of atoms.

Other notable examples

Let's look at another example using nitrogen and oxygen, which can form several oxides:

  • Nitric oxide (NO)
  • Nitrogen dioxide (NO2)
  • Dinitrogen trioxide N 2 O 3)
  • Dinitrogen tetraoxide (N 2 O 4)
  • Nitrous oxide (N2O)

Let us calculate the comparison of these compounds by taking the mass ratio:

  • In NO, 14g nitrogen combines with 16g oxygen.
  • In NO2, 14g nitrogen combines with 32g oxygen.
  • In N 2 O 3 28 grams of nitrogen combine with 48 grams of oxygen.

By focusing our attention on a fixed mass of nitrogen only (14 grams), the proportion of oxygen varies, yet still conforms to simple integer ratios that demonstrate the rule:

The oxygen mass ratio in the compounds NO, NO 2 and N 2 O 3 is 16:32:48, which reduces to 1:2:3.

Conclusion

The law of multiple proportions is not just an abstract rule, but a fundamental pillar that supports the structure of atomic theory and the complex dance of chemical reactions. It provides a systematic way to analyze and understand how elements form different types of compounds, staying true to Dalton's idea that matter is made up of atoms.

This rule extends to practical applications in laboratory scenarios and industrial processes, where understanding the stoichiometric relationships of chemical reactions is critical to efficiency, safety, and resource management.

Summary

We have learned that the law of multiple proportions shows the characteristic pattern of combination of elements, reflecting the discrete nature of atomic theory. This law reminds us that chemical unity underlies the apparent diversity of compounds, illustrating the beauty of simplicity in complexity.

The discovery of this law has deepened our understanding of chemical principles and has far-reaching implications in both theoretical and applied chemistry.


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Law of multiple proportions

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