Unit 2 Study Notes: Atomic History and Structure

Unit 2 Objectives: Atomic History and Structure

Atomic History

  • Democritus

    • Proposed that matter is made up of indivisible particles called atoms.

    • Suggested different shapes and sizes of atoms correspond to different materials.

  • Dalton

    • Formulated the atomic theory, stating:

    1. Each element is composed of extremely small particles called atoms.

    2. All atoms of a given element are identical in mass and properties.

    3. Atoms cannot be created or destroyed in a chemical reaction (conservation of mass).

    4. A chemical reaction involves reorganization of atoms to form new substances.

    • Main Flaws:

    • Atoms are not indivisible (subatomic particles exist).

    • Isotopes dispute the idea that all atoms of an element are identical in mass.

  • Rutherford

    • Conducted the gold foil experiment, which led to the discovery of the nucleus.

    • Proposed a model of the atom with a small, dense, positively charged nucleus surrounded by electrons.

  • Thomson

    • Discovered the electron using cathode rays.

    • Proposed the "plum pudding" model, where negative electrons are embedded in a positive 'soup'.

  • Bohr

    • Introduced a model where electrons travel in specific orbits around the nucleus.

    • Developed the concept of quantized energy levels for electrons.

  • Schroedinger

    • Developed quantum mechanical model of the atom, emphasizing probabilities.

    • Introduced wave functions to describe the behavior of electrons.

Atomic Structure

  • Basic Structure of the Atom

    • Atoms consist of a nucleus and subatomic particles.

    • Subatomic Particles:

    • Protons

      • Symbol: P or p

      • Charge: +1 (positive)

      • Mass: Approximately 1 amu (atomic mass unit)

    • Neutrons

      • Symbol: N or n

      • Charge: 0 (neutral)

      • Mass: Approximately 1 amu

    • Electrons

      • Symbol: e\u2212

      • Charge: -1 (negative)

      • Mass: Approximately 1/1836 amu

  • Differences between Atoms, Ions, and Isotopes

    • Atom:

    • Neutral charge, equal number of protons and electrons.

    • Ion:

    • Charged species; cation (positive charge, loss of electrons) or anion (negative charge, gain of electrons).

    • Isotope:

    • Variants of an element with the same number of protons but different numbers of neutrons.

  • Relative Masses of Subatomic Particles

    • Protons and neutrons both have a mass of approximately 1 amu.

    • Electrons have a negligible mass compared to protons and neutrons.

  • Isotopic Notation

    • Isotopic notation: A\/_Z X, where

    • A = mass number (total number of protons + neutrons)

    • Z = atomic number (number of protons)

    • X = chemical symbol of the element.

Nuclear Atom

  • Properties of Radiation

    • Alpha Radiation:

    • Composition: Helium nuclei (2 protons and 2 neutrons)

    • Relative Penetrating Power: Weak (can be stopped by paper)

    • Beta Radiation:

    • Composition: Electrons or positrons

    • Relative Penetrating Power: Moderate (can penetrate paper, but not aluminum)

    • Gamma Radiation:

    • Composition: High-energy electromagnetic radiation

    • Relative Penetrating Power: Strong (requires heavy lead or concrete to stop)

  • Concept of Half-Life

    • Definition: The time required for half of the radioactive nuclei in a sample to decay.

    • Half-Life Problem Example:

    • If a sample has a half-life of 5 years, after 10 years, one-fourth of the original sample will remain.

  • Balanced Nuclear Equations

    • Formulate equations where the sum of mass numbers and atomic numbers are balanced on both sides.

  • Nuclear Fission vs. Nuclear Fusion

    • Nuclear Fission:

    • The splitting of a heavy nucleus into two lighter nuclei, releasing energy.

    • Nuclear Fusion:

    • The combining of light nuclei to form a heavier nucleus, also releasing energy (as in stars).

Periodic Table Use/Chemical Nomenclature

  • Key Definitions on the Periodic Table

    • Atomic Number: Number of protons in an atom.

    • Mass Number: Total number of protons and neutrons.

    • Average Atomic Mass: Weighted average of all isotopes of an element.

  • Calculating Average Atomic Mass

    • Use:
      ext{Average Atomic Mass} = rac{ ext{Sum of (Isotope Mass} imes ext{Fractional Abundance)}}{ ext{Total of Abundances}}

  • Naming and Symbol Writing Rules

    • Atoms are represented by their elemental symbols.

    • Cations (positive ions) include Roman numerals to indicate charge (e.g., Iron(III) for Fe\^3+).

    • Anion names end in –ide (e.g., chloride for Cl\^-).

  • Using the Periodic Table

    • Determine number of protons, neutrons, electrons, atomic numbers and mass numbers based on element information.

    • For ions, adjust based on the charge of the ion.

  • Isotopic Notation Reading

    • Understand how to read different types of isotopic notation provided in the periodic table.

The Mole

  • Definition of the Mole

    • A mole is defined as the quantity of substance that contains as many entities (atoms, molecules, etc.) as there are atoms in 12 grams of carbon-12 (approximately 6.022 \times 10^{23} entities).

  • Calculating Molar Mass

    • Molar mass is the mass of one mole of a substance, usually expressed in grams per mole (g/mol).

  • Dimensional Analysis in Moles

    • Method to convert between mass, moles, and number of particles:

    • Example Calculation:
      ext{mass (g)} \rightarrow ext{moles} \rightarrow ext{number of particles}

    • Use conversion factors based on molar mass to facilitate the conversion.