Atomic Theory

The Atom in Modern Chemistry

  • Chapters 1.1 – 1.6

    • Key Topics:

      • Matter: properties and classification

      • Early Atomic Theory

      • J.J. Thomson

      • The Nuclear Model of the Atom

      • Atomic Number

      • Ions

      • Atomic Mass and Isotopes

      • Avogadro’s Number and the Mole

Adaptation of Material

  • Slides are adapted from Dr. Frank Ow and Oxtoby Textbook.

Matter and Its Properties and Transformations

  • Definition of Matter:

    • Matter is any substance that has mass and occupies space.

    • Fundamental building blocks of matter are atoms, which cannot be chemically broken down.

    • Atoms combine to form larger molecules and compounds.

    • States of matter include:

      • Gases

      • Solids

      • Liquids

Classification of Matter

  • Decision Flow:

    • Is it pure?

      • Yes:

        • Pure Substances

        • Can be classified as:

          • Element: Basic, cannot be decomposed (e.g., Na, Zn).

          • Compound: Combinations of elements (e.g., H₂O, NaCl).

      • No:

        • Mixture:

          • Homogeneous (Solutions): Uniform composition.

          • Heterogeneous: Non-uniform composition, multiple phases.

Pure Substances

  • Elements: Atomic units that cannot be broken down (e.g., Hydrogen, Oxygen).

  • Compounds: Combinations of different elements that can be decomposed.

  • Example: Water (H₂O) comprises hydrogen and oxygen atoms.

Dalton’s Atomic Theory (1808)

  • Atoms are the smallest identifiable units of matter and are indestructible.

  • All atoms of a given element are identical, distinct from those of other elements.

  • Atoms combine in small whole-number ratios to form compounds.

Law of Conservation of Mass (1785, Lavoisier)

  • Mass is conserved in chemical reactions; it cannot be created or destroyed.

Law of Multiple Proportions

  • When elements form compounds, the ratios of the masses of the first element to a fixed mass of the second are small whole numbers.

  • Example: Water (H₂O) forms from 2.02 g of hydrogen and 16.00 g of oxygen, while hydrogen peroxide (H₂O₂) uses 2.02 g of hydrogen with 32.00 g of oxygen.

Law of Definite Proportions

  • In any chemical compound, the mass ratios of the elements in that compound are fixed, regardless of the source.

  • Example: Water's mass ratio is always 1.01 g of hydrogen for every 8.00 g of oxygen (11.2% H and 88.8% O).

J.J. Thomson and the Electron

  • Thomson's Discoveries (1897): Identified negatively charged electrons.

  • Millikan’s oil drop experiment (1909) quantified the charge and mass of an electron.

Thomson’s Model of the Atom (Plum-Pudding Model)

  • Proposed that an atom consists of a positively charged sphere with electrons dispersed throughout.

Rutherford’s Gold Foil Experiment (1911)

  • Discovered that atoms have a dense, positively charged nucleus, containing most of the atom's mass.

Nuclear Model of the Atom

  • Structure:

    • Nucleus: Contains protons and neutrons.

    • Electrons orbiting the nucleus.

Atomic Number

  • Defined as the number of protons in an atom's nucleus; also equals the number of electrons in neutral atoms.

  • Periodic table is arranged by increasing atomic number.

Ions

  • Atoms can gain or lose electrons to form charged ions:

    • Cations: Positively charged (lose electrons).

      • Example: Na → Na⁺ + e⁻

    • Anions: Negatively charged (gain electrons).

      • Example: O + 2e⁻ → O²⁻

Atomic Mass and Isotopes

  • Isotopes: Atoms of the same element with different neutron numbers.

  • Atomic mass number = number of neutrons + number of protons.

  • Examples in Carbon: C-12, C-13, C-14.

  • Carbon's isotopes distribution: 98.93% C-12, 1.07% C-13.

Atomic Mass Calculation

  • Average atomic mass accounts for the weighted abundance of all isotopes.

  • Example: Chlorine isotope calculation using relative abundance and exact mass.

The Mole

  • Definition: Represents the quantity of atoms in 12 g of C-12.

  • Avogadro’s Number: 1 mole = 6.022 × 10²³ units.

  • Conversion: To convert to/from moles, use Avogadro’s number.

Molar Mass of Atoms

  • Molar mass (g/mol) equals the atomic mass in amu.

  • To convert units:

    • Grams to moles: divide by molar mass.

    • Moles to grams: multiply by molar mass.

Molar Mass of Compounds

  • Calculated by summing the atomic masses of constituent atoms.

  • Example calculations:

    • CO₂: 44.01 g/mol

    • CaCO₃: 100.09 g/mol

    • C₈H₁₈: 114.26 g/mol.

Using Avogadro’s Number for Compounds

  • Applied to find the number of molecules in a sample based on its mass and molar mass.