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 definite proportions


The law of definite proportions, also called the law of constant composition, is a fundamental principle in chemistry first articulated by French chemist Joseph Proust in 1797. This law states that a chemical compound will always contain its constituent elements in a fixed proportion by mass, regardless of the source or size of the sample. Simply put, every pure sample of a compound is composed of the same elements in the same proportions by weight.

Understanding the law

To understand this law better, let's understand it with an example. Consider the compound water (H2O). Water is composed of hydrogen and oxygen. According to the law of definite proportions, every sample of pure water will always have the same ratio of hydrogen and oxygen, which is approximately 2:16 or 1:8 by mass. This means that in any given volume of water, the mass of oxygen will always be eight times the mass of hydrogen.

H2O: 
    2 parts hydrogen 
+ 16 parts oxygen,
  = 18 parts of H2O

No matter where you take a water sample from, whether it is from a river, sea, rainwater or distilled in a laboratory, if it is pure water, the mass ratio of hydrogen and oxygen will always be 1:8. This constancy is what the law of definite proportions guarantees.

Theoretical basis and significance

The law of definite proportions is important because it laid the groundwork for chemical formulas and the concept of stoichiometry in chemistry. Prior to Joseph Proust's research, there was debate among chemists as to whether elements could combine in any proportions. Proust's experiments demonstrated that specific compounds are bound by definite mass ratios.

This discovery played a key role in the development of atomic theory. It confirmed that matter is made up of atoms, which combine in specific ways to form compounds. If atoms are tiny particles with uniform masses, it makes sense that they would combine in fixed, whole number ratios. This insight led to the chemical formulas we use today. A formula like CO2 reflects the specific ratio of carbon and oxygen atoms that combine to form carbon dioxide.

Examples of fixed ratios

Sodium chloride

Sodium chloride, commonly known as table salt, is another example. The chemical formula for sodium chloride is NaCl, indicating that the sodium (Na) and chlorine (Cl) atoms are in a 1:1 ratio, which translates to a mass ratio of about 23:35.5. Therefore, in every sample of sodium chloride, the ratio of the mass of sodium to the mass of chlorine will always be about 23:35.5.

Sodium Chloride: 
    23 parts sodium 
+ 35.5 parts chlorine,
   = 58.5 parts NaCl

No matter where you get the salt from, this mass ratio remains constant as long as it is pure sodium chloride.

Carbon dioxide

Carbon dioxide, widely known as a greenhouse gas, also follows this rule. Its chemical formula is CO2, which specifies the ratio of one carbon atom to two oxygen atoms. By mass, this is equivalent to about 12 parts carbon and 32 parts oxygen, for a total of 44 parts. Thus, each time carbon dioxide is formed, its mass ratio remains constant.

CO2: 
    12 parts carbon 
+ 32 parts oxygen,
  = 44 parts CO2

Visualizing ratios with a simple analogy

To understand this concept better, let's use a simple example of a fruit salad. Imagine that you use exactly two parts apples and one part bananas to make a fruit salad. No matter how small or large the salad is, as long as the ratio of apples and bananas remains constant, the taste and texture of the salad does not change. Similarly, the law of definite proportions ensures that no matter how much a chemical compound is used, the elemental composition remains the same.

SVG visualization example: water molecule structure

Consider the illustration of the water molecule below, which shows hydrogen and oxygen atoms in a stable 2:1 ratio.


    
    
    
    
    
    O
    H
    H

Challenges and exceptions to the law

Though the law of definite proportions is widely applicable, there are certain exceptions and challenges associated with it:

  • Non-stoichiometric compounds: Some compounds, especially metal oxides, can have variable composition. For example, iron oxide can exist as FeO, Fe2O3 etc., depending on the method and conditions of its formation.
  • Isotopic variations: Different isotopes of an element may cause the mass ratio to change even though the atomic ratio remains constant.
  • Impurities: In practical scenarios, impurities can affect the observed mass structure unless the compound is purified.

Conclusion

The law of definite proportions is a cornerstone of chemical science, essential to understanding chemical reactions and compound formation. It confirms that the elemental composition of any chemical substance is consistent and predictable, which contributes to our ability to use chemical formulas to succinctly describe compounds. This law not only supports atomic theory, but also provides consistency and reliability in chemical analysis and synthesis, proving indispensable in scientific discoveries and industrial applications alike.

Although exceptions to this rule may exist due to non-stoichiometric compounds and isotopic variations, the principle remains central to classical chemistry and its teaching today, and serves as a bridge to advanced understanding of molecular chemistry and beyond.


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Law of definite proportions

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