Ancient Indian and Greek Philosophers’ Views on Matter:

  • Indian philosophers (around 500 BC), like Maharishi Kanad, believed that matter can be divided into smaller and smaller particles. Eventually, there would be a point where further division is not possible, and these smallest particles were called Parmanu.
  • Another philosopher, Pakudha Katyayama, expanded on this idea, suggesting that these tiny particles often combine to form various kinds of matter.
  • Greek philosophers, like Democritus and Leucippus, also proposed that matter can be divided until the indivisible particles, which they called atoms (from the Greek word for indivisible).
  • These ideas, though early, were philosophical rather than experimental.

Laws of Chemical Combination:

By the 18th century, scientists began to differentiate between elements and compounds. Antoine Lavoisier and Joseph Proust developed two important laws of chemical combination through experimentation:

  1. Law of Conservation of Mass (Lavoisier):
    • Mass is neither created nor destroyed in a chemical reaction. The total mass before and after a reaction remains the same.
    • Example: In a chemical reaction, when a substance reacts with another, the total mass of the products equals the total mass of the reactants.
  2. Law of Constant Proportions (Proust):
    • In a given compound, the elements always combine in a fixed, constant ratio by mass, no matter the source or method.
    • Example: In water, the ratio of hydrogen to oxygen by mass is always 1:8, regardless of where the water is obtained.

John Dalton’s Atomic Theory:

  • Dalton’s atomic theory (1808) helped explain the Law of Conservation of Mass and Law of Constant Proportions. His theory stated that:
    1. All matter is made of tiny particles called atoms.
    2. Atoms cannot be created or destroyed in a chemical reaction.
    3. Atoms of the same element are identical in mass and properties.
    4. Atoms of different elements have different masses and properties.
    5. Atoms combine in simple whole-number ratios to form compounds.
    6. The relative number and types of atoms are constant in a given compound.
  • Dalton’s theory also emphasized the atomic mass of elements and helped to clarify the concept of chemical reactions.

Atoms and Their Properties:

  • Atoms are extremely small, with radii measured in nanometers (nm). For comparison, the size of an atom is around 10-10 m, while the size of a grain of sand is much larger (10-4 m).
  • Modern symbols for elements, like H for hydrogen and O for oxygen, are standardized by the IUPAC (International Union of Pure and Applied Chemistry).

Atomic Mass:

  • Dalton’s atomic theory introduced the idea that each element has a specific atomic mass.
  • The atomic mass unit (amu), later standardized as unified atomic mass (u), was initially based on oxygen’s mass. Later, the carbon-12 isotope became the standard reference for atomic mass.

Key Insights and Competetive Exam Insights:

  1. Atoms’ Indivisibility:
    • Atoms are indivisible in Dalton’s theory, which aligns with the Greek view, but later studies revealed that atoms themselves are made up of subatomic particles.
  2. Scientific Concepts:
    • Law of Conservation of Mass is fundamental in chemical reactions and is a cornerstone of modern chemistry.
    • Law of Constant Proportions explains why compounds, like water (H2O), always have elements in fixed ratios (hydrogen and oxygen in 1:8 by mass).
  3. Applications:
    • The atomic theory and laws of chemical combination help explain many phenomena, including chemical reactions and the formation of compounds, which is important for understanding chemical processes in industry, medicine, and natural sciences.

Practical Applications and Understanding:

  • Experimental Insights: In experiments, like the reaction between sodium carbonate and acetic acid, we can apply the Law of Conservation of Mass to verify that the total mass of reactants equals the mass of products after the reaction. This experiment emphasizes the importance of measuring before and after a reaction.

How Do Atoms Exist?

  • Atoms of most elements cannot exist alone. Instead, they combine to form molecules and ions. These molecules or ions then group together to form matter, which we can see, touch, or feel.

Atomic Mass Unit

  • The atomic mass unit (AMU) is a standard unit of mass used to express atomic and molecular masses. It’s defined as 1/12 of the mass of a carbon-12 atom.

Why Can’t Atoms Be Seen?

  • Atoms are incredibly small, much smaller than the wavelength of visible light. Therefore, they cannot be seen by the naked eye.

Molecule: Definition

  • A molecule is a group of two or more atoms chemically bonded together. It is the smallest particle of a substance that can exist independently and still show all the properties of that substance.
  • Example: The molecule of oxygen (O₂) consists of two oxygen atoms.

Types of Molecules

  1. Molecules of Elements:
    • Consist of atoms of the same element (e.g., O₂ is a diatomic molecule).
    • Atomicity: The number of atoms in a molecule.
      • Monoatomic: Argon (Ar), Helium (He).
      • Diatomic: Oxygen (O₂), Hydrogen (H₂), Nitrogen (N₂), etc.
      • Polyatomic: Phosphorus (P₄), Sulfur (S₈).
  2. Molecules of Compounds:
    • Made from atoms of different elements in fixed proportions.
    • Example: Water (H₂O) is made from 2 hydrogen atoms and 1 oxygen atom.

Atomic Mass and Molecules

  • For calculating the number of atoms in a compound, we use atomic mass units (AMU).
    • Example: In water (H₂O), hydrogen has an atomic mass of 1, and oxygen has an atomic mass of 16. The molecular mass of H₂O = (2 × 1) + (1 × 16) = 18 AMU.

Ions and Their Types

  • Ions: Atoms or groups of atoms that carry an electric charge.
    • Anion: Negatively charged ion (e.g., chloride ion Cl⁻).
    • Cation: Positively charged ion (e.g., sodium ion Na⁺).
  • Polyatomic Ions: Groups of atoms that together carry a charge (e.g., ammonium NH₄⁺).

Formation of Ions

  • Example: In sodium chloride (NaCl), sodium (Na) loses an electron to become Na⁺ (cation), and chlorine (Cl) gains an electron to become Cl⁻ (anion).

Valency

  • Valency is the combining power of an element, or how many bonds an atom can form with other atoms.
    • Example: Sodium (Na) has a valency of 1 (forms 1 bond), while oxygen (O) has a valency of 2 (forms 2 bonds).
  • Valency is like the number of “arms” an atom has to bond with other atoms.

Writing Chemical Formulae

  • Chemical Formula: A symbolic way to represent the composition of a compound.
    • Example: The formula of water is H₂O (2 hydrogens and 1 oxygen).
  • Ionic Compounds:
    • For ionic compounds, the formula is written by balancing the charges of the ions.
    • Example: Magnesium chloride (MgCl₂) – Magnesium (Mg²⁺) and Chlorine (Cl⁻) balance to form the formula MgCl₂.
  • Rules for Writing Formulae:
    1. Metal first, then non-metal (e.g., NaCl, CaO).
    2. For polyatomic ions, use brackets if there’s more than one ion (e.g., Na₂SO₄).

Examples:

  • Sodium Oxide: Na₂O
  • Aluminium Chloride: AlCl₃
  • Magnesium Hydroxide: Mg(OH)₂

Molecular Mass and Formula Unit Mass

  1. Molecular Mass: The sum of the atomic masses of all atoms in a molecule.
    • Example: For H₂O, the molecular mass = 18 AMU (2 × 1 for hydrogen + 16 for oxygen).
  2. Formula Unit Mass: Similar to molecular mass, but used for ionic compounds. It’s the sum of the atomic masses in the simplest formula unit of a compound.
    • Example: For NaCl, the formula unit mass = 23 (Na) + 35.5 (Cl) = 58.5 AMU.

THESE ALL ARE THE NOTES OF CHAPTER 3 SCIENCE. AND AFTER SOME TIME YOU GET IMPORTANT QUESTIONS HERE. *#THANKS FOR VISITING, VISIT AGAIN#* 😊