Chapter 2

Atomic Theory and Subatomic Particles (Sections 2.1 & 2.2)

• Evolution of Atomic Theory

  • The atomic theory has undergone significant changes since its inception by John Dalton.

  • Dalton proposed that all matter is composed of atoms, indivisible and indestructible particles, unique to each element.

  • Later discoveries involving subatomic particles (electrons, protons, and neutrons) refined this model substantially.

• Use of Ratios in Compounds

  • Ratios of elements are crucial for distinguishing between identical and different compounds.

  • For instance, water (H₂O) and hydrogen peroxide (H₂O₂) have the same elements but different ratios, leading to distinct properties.

• Key Experiments in Understanding Subatomic Particles

  • Various groundbreaking experiments have shaped our understanding of atomic structure:

  • Thomson’s Cathode Ray Experiment: Established the existence of electrons and led to the 'plum pudding' model of the atom, suggesting electrons are dispersed within a positively charged 'soup'.

  • Rutherford’s Gold Foil Experiment: Disproved the plum pudding model by showing that atoms have a small, dense nucleus, surrounded by electrons.

  • Chadwick’s Discovery of Neutrons: Added to the atomic model by identifying neutrons, which have no charge and contribute to atomic mass.

• Models of the Atom

  • The results from these experiments have led to the development of various atomic models, including:

  • Dalton's model (solid indivisible sphere)

  • Thomson's plum pudding model

  • Rutherford's nuclear model

  • Bohr's model (quantized energy levels)

Atomic Structure, Symbols, and Calculations (Sections 2.3 & 2.4)

• Location and Properties of Subatomic Particles

  • Protons: Found in the nucleus, positively charged, with a mass of approximately 1 atomic mass unit (amu).

  • Neutrons: Also located in the nucleus, neutral charge, similar mass to protons (1 amu).

  • Electrons: Orbit the nucleus, negatively charged, much lighter (~1/1836 amu) compared to protons and neutrons.

• Isotope Symbols (look at notes)

  • Isotope notation can be represented as:

  • C or (^{13}C), where the superscript denotes the mass number (protons + neutrons) and the subscript denotes the atomic number (number of protons).

  • Cations and Anions:

  • Cation: positively charged ion (loss of electrons) e.g., (Na^+).

  • Anion: negatively charged ion (gain of electrons) e.g., (Cl^-).

• Determining Subatomic Counts

  • For a neutral atom:

  • Number of protons = Number of electrons

  • For ions:

  • Subtract or add electrons according to the charge to obtain the number of electrons.

  • Count of neutrons is determined by (\text{Mass Number} - \text{Atomic Number} = \text{Number of Neutrons}).

• Definitions of Key Terms

  • Atomic Number (Z): Number of protons in the nucleus of an atom, unique to each element.

  • Mass Number (A): Total number of protons and neutrons in an atom.

  • Atomic Mass Unit (amu): A standardized unit of mass that quantifies mass on an atomic or molecular scale (1 amu is defined as one twelfth of the mass of a carbon-12 atom).

  • Average Atomic Mass: Weighted average of the masses of all isotopes of an element, reflecting their natural abundance.

  • Molecular Mass: Sum of atomic masses of the atoms in a molecule.

  • Molar Mass: Mass in grams of one mole of a substance, numerically equal to its average atomic mass in amu.

• Using Isotopic Abundance ( look at notes)

  • Isotopic abundance indicates the relative amount of each isotope of an element in nature, which can be used to calculate the average atomic mass (
    (\text{Average Atomic Mass} = \sum (\text{Isotopic Mass}) \times (\text{Fractional Abundance}))).

  • Alternatively, average atomic mass can help estimate the abundance of isotopes when one or the other is known.