History of the Atom and Early Theories

Discovery and History of the Atom

• Maharishi Kanada, an Indian philosopher, discovered the concept of the atom in ancient India, stating matter is composed of the smallest unit, "Parmanu".
• Democritus and Lucippus, ancient Greek philosophers, proposed that all matter is made of small, indivisible particles called atoms.

Early Atomic Theory

• Antoine Lavoisier provided evidence that atoms are the building blocks of matter.
• Lavoisier described atoms as very small and spherical, with a radius around 1 nanometer (10⁻⁹ meters).
• Lavoisier believed atoms were indivisible.

Laws of Chemical Combination

• Antoine Lavoisier proposed the Law of Conservation of Mass and the Law of Constant Proportions.
  • Law of Conservation of Mass: In a chemical reaction, the total mass of reactants equals the total mass of products.
  • Law of Constant Proportions: A pure chemical compound always contains the same elements in the same proportions by mass.

### Examining the Law of Conservation of Mass

• Reactants are substances before a chemical reaction, and products are substances after the reaction.
• The Law of Conservation of Mass states that the total mass of reactants equals the total mass of products.
• If 20 g of reactant X reacts with 20 g of reactant Y to produce a new compound, the new compound will have a mass of 40 g.
• Melting ice (a physical change) also demonstrates this law: 20 g of ice will produce 20 g of water.

Examining the Law of Constant Proportions

• Water (H₂O) always contains hydrogen and oxygen in a fixed 1:8 mass ratio.
• This constant ratio ensures the consistent formation of water from hydrogen and oxygen. A different ratio (e.g., 3:8) will not result in water.

Dalton's Atomic Theory

• John Dalton's atomic theory, a cornerstone of modern chemistry, was developed in the 19th century.
• Dalton's postulates:
  • All matter is made of tiny, indivisible particles called atoms.
  • Atoms of the same element are identical in mass and properties, while atoms of different elements have different masses and properties.
  • Atoms cannot be created or destroyed, only rearranged.
  • When elements react, atoms combine in simple whole number ratios to form compounds.
• Dalton's concept of indivisibility was later disproven. Subatomic particles (protons, neutrons, electrons) were discovered.

Dalton's Atomic Theory

• Atoms are the smallest particles retaining an element's properties.
• Atoms are indivisible and indestructible.
• Atoms of the same element have the same mass and properties; atoms of different elements differ in mass and properties.
• Atoms of different elements combine in whole-number ratios to form compounds.
• Compounds have a fixed proportion of atoms (Law of Constant Proportions).

Drawbacks of Dalton's Atomic Theory

• Atoms are divisible, due to subatomic particles like protons, neutrons, and electrons.
• Atoms of the same element can have different masses (isotopes).
• Different elements can have the same mass (isobars).

Modern Symbols for Elements

• Scientists developed a simplified system, using elements' names to represent them.
• One-letter symbols are derived from the first letter (e.g., H for Hydrogen).
• Two-letter symbols use the first two letters (e.g., He for Helium).
• Some symbols are derived from Latin names (e.g., Na for Sodium, from Natrium).

Mnemonic to Remember Elements

• First 10: "Hi, Hi, Hello! Hi, Hello! BBC News On Friday Night!"
• Next 11: "Na (No) Mangoo Abdool Se, Pepsi Soda Cold Drink, Aur (And) Kaaju Katli!"

Atomic Mass

• Atomic mass is the mass of an atom.
• IUPAC determines atomic masses.
• The standard is the Carbon-12 atom (C-12). One atomic mass unit (amu or u) is 1/12th the mass of a C-12 atom.

Calculating Atomic Mass

• Even-numbered elements: Atomic mass ≈ element's position × 2.
• Odd-numbered elements: Atomic mass ≈ element's position × 2 + 1.
• Exceptions: Hydrogen (1 u), Beryllium (9 u), Nitrogen (14 u), Argon (40 u).

Atomic Existence

• Atoms combine to achieve stability.
• Atoms exist as:
  • Single atoms (e.g., noble gases, hydrogen).
  • Diatomic molecules (e.g., H₂, O₂, Cl₂).
  • Polyatomic molecules (e.g., P₄, S₈).

Ions

• Ions are charged atoms or groups of atoms.
• Formed by gaining or losing electrons.
• Losing electrons results in a positive ion (cation); gaining electrons, a negative ion (anion).

Introduction to Ions

• Atoms become ions when gaining or losing electrons.
• A neutral atom forms a cation (positive ion) by losing electrons, and an anion (negative ion) by gaining electrons.

Types of Ions

• Simple ions contain a single element (e.g., Na⁺, H⁺, Cl⁻).
• Compound ions (polyatomic ions) contain multiple elements (e.g., NH₄⁺, OH⁻, SO₄²⁻).

Concept of Atomicity

• Atomicity is the number of atoms in a molecule of an element or compound.

Types of Atomicity

• Monoatomic (one atom).
• Diatomic (two atoms).
• Triatomic (three atoms).
• Polyatomic (more than three atoms).

Calculating Atomicity

• Determine atomicity by counting atoms in the chemical formula.

Formula Unit Mass

• Formula unit mass is the sum of atomic masses in a compound's formula unit. (e.g., H₂O: 18 u).

Molecular Mass

• Molecular mass is the sum of atomic masses of all atoms in a molecule. (e.g., H₂O: 18 u).

How to Calculate Molecular Mass

• Use the Periodic Table for atomic masses.
• Multiply each element's atomic mass by its subscript in the formula.
• Add the results for all elements in the compound. (e.g., CaCO₃: 100.1 u).

Importance of Ions

• Ions are crucial for chemical bonding and reactions.
• Understanding ionic charges is essential for naming and writing chemical formulas.

Remembering the Periodic Table

• Mnemonic: "Na Kaho Agar Cute Tujhe, Maange Car Jan, Ferrari, Kyu Nahin".
• Valences for elements and some polyatomic ions are represented in the mnemonic.

Atoms Combining

• Atoms combine for stability by filling their valence shells.
• Valence reflects an element's bonding capacity.

Naming Compounds

• Use valences to name and formulate compounds.
• Valences are used to create a chemical formula.

Important Concepts

• Law of Conservation of Mass: Reactant mass = Product mass.
• Mole: 1 mole = 6.022 × 10²³ (Avogadro's number).
• Molar Mass: The mass of 1 mole of a substance. Convert atomic mass units to grams.
• Moles calculations: Moles = Mass/Molar mass; Moles = Number of particles/Avogadro's number.

Top 5 Important Questions

• Question 1: 10 g Silver Nitrate and 10 g Sodium Chloride react. What happens to the mass after the reaction?
  • Answer: Mass remains the same. The Law of Conservation of Mass applies.