Atoms, Molecules, and Ions

• Chapter 2: Atoms, Molecules, and Ions

• Atomic Theory of Matter

• Greek philosophers like Democritus proposed that the smallest particle of matter is the atom ("atomos" meaning uncuttable).

• John Dalton developed a formal atomic theory in 1803:

  • Law of Constant Composition: A compound always contains the same proportions of elements by mass.

  • Law of Conservation of Mass: Mass is neither created nor destroyed in chemical reactions.

  • Law of Multiple Proportions: If two elements can form more than one compound, the different masses of one element that combine with a fixed mass of the other are in small whole number ratios.

• Postulates of Dalton’s Atomic Theory (1808):

1. All matter is made of atoms, which are indivisible.

2. All atoms of a given element are identical, while atoms of different elements differ.

3. Compounds contain atoms of more than one element in fixed ratios.

4. Chemical reactions involve rearrangements of atoms; no atoms are created or destroyed.

• Law of Conservation of Mass

• The total mass before a chemical reaction is equal to the total mass after. For example, in the reaction 3H2 + N2 → 2NH3, the mass of reactants equals the mass of products.

• Discovery of Subatomic Particles:

• Initially, atoms were thought to be indivisible; later discoveries revealed smaller particles:

  • Electrons: Negative particles discovered in cathode rays by J.J. Thomson (1897).

  • Millikan’s Oil-Drop Experiment (1909): Determined the charge of an electron at approximately 1.602 x 10^-19 coulombs.

  • Radioactivity: Discovered by Henri Becquerel, later studied by the Curies; showed that atoms can emit high-energy radiation, indicating subatomic structures.

  • Nucleus: Ernest Rutherford's gold foil experiment led to the understanding that atoms have a dense, positively charged nucleus surrounded by electrons.

• Subatomic Particles:

• Protons (p+): Positive charge, mass ≈ 1 amu; located in the nucleus.

• Neutrons (n°): No charge, mass ≈ 1 amu; located in the nucleus.

• Electrons (e-): Negative charge, negligible mass compared to protons/neutrons; found outside the nucleus.

• Atomic Number and Mass Number:

Atomic Number (Z): Number of protons in an atom's nucleus (equals the number of kelectrons in a neutral atom).

• Mass Number (A): Total number of protons and neutrons in the nucleus.

• Isotopes:

• Atoms of the same element having the same atomic number but different mass numbers (due to varying numbers of neutrons). Example: Carbon isotopes include Protium (C-12), Deuterium (C-13), and Tritium (C-14).

• Atomic Mass Unit (amu):

• Small mass unit used for atoms, where 1 amu = 1.66054 x 10^-24 g.

• Atomic weights are averages based on isotopic abundance, compared to C-12, defined as exactly 12 amu.

• The Periodic Table:

• Organized by periods (rows) and groups (columns); elements are categorized based on similar properties.

• Group Designations:

  • Group 1: Alkali metals

  • Group 2: Alkaline earth metals

  • Group 17: Halogens

  • Group 18: Noble gases

• Types of Elements:

• Metals: Good conductors, malleable, ductile.

• Nonmetals: Poor conductors, brittle, can be gases.

• Metalloids: Intermediate properties, can behave as metals or nonmetals.

• Chemical Formulas:

• Indicate the number and type of atoms in a molecule.

• Ions: Charged atoms that can be cations (positive) or anions (negative).

• Naming Compounds:

• Ionic Compounds: Typically formed from metals (cations) and nonmetals (anions). Cation named first, anion receives “-ide”.

• Molecular Compounds: Consist of nonmetals; prefixes indicate the number of atoms (mono- for 1, di- for 2, etc.).

• Acids and Bases:

• Acids yield H+ in solution (e.g., HCl).

• Bases yield OH- in solution (e.g., NaOH).

• Hydrates are compounds incorporating water molecules (e.g., CuSO4·5H2O).

• Common Names and Systematic Names of Compounds:

• Familiar names versus systematic IUPAC names (e.g., H2O is water but systematically referred to as dihydrogen monoxide).

• End of Chapter 2