Unit2 (1)

Atomic Structure Notes

A. Early Chemical Laws

  • Law of Conservation of Mass

    • Proposed by Lavoisier (1780's)

    • States that the total mass of reactants equals the total mass of products during a chemical reaction.

  • Laws of Definite and Multiple Proportions

    • Proposed by Dalton (1800's)

    • Definite Proportions: Compounds are always made from the same elements in the same proportions.

    • Multiple Proportions: Two elements can combine in different ratios to form more than two different compounds. For example, CO, H2O, NO, CH4, C2H6, CO2, H2O2, and NO2.

    • Ratios of element A's mass combining with a constant mass of element B reduce to whole numbers. For instance, sulfur (S) combines with oxygen (O) in different ratios:

      • X = SO3 (Sulfur Trioxide) and Y = SO2 (Sulfur Dioxide).

B. Work of Joseph Gay-Lussac

  • At the same temperature and pressure, equal volumes of gases contain equal numbers of particles.

  • Avogadro's Law states that equal volumes of gases at the same temperature and pressure contain equal numbers of molecules.

C. Dalton's Atomic Theory (1808)

  1. Elements are composed of indivisible atoms.

  2. Atoms of the same element have identical properties, while atoms of different elements differ in properties.

  3. Compounds are formed when atoms combine in fixed proportions.

  4. Atoms do not change their identity during reactions; they simply recombine.

D. Modern Atomic Structure

  • Atoms consist of subatomic particles:

    • Electrons: Discovered by Thomson (1899).

      • Cathode rays are streams of electrons observed in cathode ray tubes.

      • Measured the mass-to-charge ratio of electrons.

      • Concluded that all atoms contained negatively charged electrons, implying the existence of positively charged particles in the atom.

E. Important Discoveries in Atomic Physics

  1. Millikan's Oil Drop Experiment (1909)

    • Determined the charge of the electron as approximately 1.6 x 10^-19 coulombs.

    • Mass of an electron is approximately 9.11 x 10^-28g.

  2. Rutherford's Gold Foil Experiment (1911)

    • Shot alpha particles at a thin foil of gold.

    • Most particles passed through, but some were deflected, revealing a dense, positively charged nucleus within the atom.

    • Established that the atom is mostly empty space and that the nucleus contains protons and neutrons.

F. Basic Properties of Atoms and Isotopes

  • Atomic Mass and Isotopes:

    • Atomic mass number is the sum of protons and neutrons in the nucleus.

    • Isotopes are atoms of the same element with different numbers of neutrons; e.g., 12C (normal carbon) and 14C (radioactive carbon).

  • Atomic Mass Unit (amu) defined as 1/12 the mass of a 12C atom ( ~1.66 x 10^-24g).

G. Atomic Particles and Their Masses

  • Proton:

    • Charge: +1.00 e

    • Mass: 1.0073 amu

  • Neutron:

    • Charge: 0

    • Mass: 1.0087 amu

  • Electron:

    • Charge: -1.00 e

    • Mass: 9.11 x 10^-28g

H. Percent Composition and Empirical Formula

  • Mass Percent Composition: Utilized to understand the contribution of each element in a compound.

    • Use to find the empirical formula: the simplest whole number ratio of atoms.

    • Example: Analyze compounds X and Y to find their empirical formulas (e.g., X found to be NO2).

I. Molar Mass and Calculation Example

  1. The molar mass can be derived from the atomic masses of the constituent atoms.

  2. Sample calculation showing how to deduce the formula and molar mass based on mass percentages provided.

J. Practical Applications of Molar Mass

  • Avogadro's Number (NA): Represents the number of atoms, molecules, or ions in one mole (6.02 x 10^23).

  • Used for calculating the number of particles from grams of a substance, enabling calculations for compounds like CO2.