ch 3

Introductory Chemistry - Chapter 3 Notes

History of Atomic Theory

  • Democritus (400 B.C.E.)
    • Proposed that matter is made of indivisible particles termed "atomos."

Law of Conservation of Mass

  • Matter cannot be created or destroyed during a chemical reaction.
  • Example of Equation:
    • Hydrogen + Oxygen ( 4.0 ext{ g} + 32.0 ext{ g} = 36.0 ext{ g Water} )
  • Antoine Lavoisier (1743-1794) formalized the law stating that the total mass of reactants equals the total mass of products.
  • Example:
    • Reaction involving methane and oxygen:
    • ( 16.0 ext{ g of CH}4 + 64.0 ext{ g of O}2 = 44.0 ext{ g of CO}2 + 36.0 ext{ g of H}2O )
    • Calculate the mass of carbon dioxide produced:
      • ( 16.0 + 64.0 = 80.0 \text{ g total reactants} )
      • ( 80.0 - 36.0 = 44.0 \text{ g of CO}_2 )

Foundations of Atomic Theory

  • John Dalton (1766-1844):
    • Proposed that elements consist of small, indivisible atoms.
    • Each type of atom is unique and combines in whole-number ratios to form compounds.
    • Atoms remain unchanged during chemical reactions.

Fundamental Concepts of Atomic Theory

  1. All matter comprises atoms.
  2. Each element's atoms have distinct characteristics.
  3. Atoms are unaltered in reactions but combine in fixed ratios to form compounds.

Visualization of Atoms

  • X-ray Crystallography:
    • A technique used by scientists to visualize atomic arrangements.
    • Example Reference: PDB ID: 1GZX, Paoli et al, 2002.

Periodic Table of Elements

  • Historical Context: Developed to organize elements based on atomic structure.
  • Mendeleev's Contribution:
    • Created the first periodic table which organized elements by increasing atomic mass and grouped elements by similar properties.
  • Structure: Elements displayed in rows called periods and columns called groups/families.

Element Classification

  • Main Groups: Elements are categorized based on their properties. They include Metals, Metalloids, and Nonmetals.
  • Location on the Periodic Table:
    • Metals are typically on the left;
    • Nonmetals are located in the upper right;
    • Metalloids are in between both categories.

Detailed Group Characteristics

  • Group 1A: Alkali Metals
    • Characteristics: Soft metals, react violently with water.
  • Group 2A: Alkaline Earth Metals
    • Characteristics: Less reactive than alkali metals, burn brightly when ignited.
  • Group 7A: Halogens
    • Characteristics: Exist as diatomic molecules, form various compounds.
  • Group 8A: Noble Gases
    • Characteristics: Generally inert, do not readily form compounds, gaseous at room temperature.

Atomic Structure

  • Atoms consist of subatomic particles: protons, neutrons, and electrons.
  • Subatomic Particles Characteristics:
    • Proton: Mass of (1.0073 ext{ u}), charge of +1;
    • Neutron: Mass of (1.0087 ext{ u}), charge of 0;
    • Electron: Mass of (0.0005 ext{ u}), charge of -1.

Key Concepts in Atomic Structure

  • Atomic Identity: Determined by the number of protons, affecting the element type.
  • Atomic Number: Number of protons in an atom, equals the number of electrons in a neutral atom.
  • Mass Number: Total number of protons and neutrons in an atom.
  • Isotopes: Atoms with the same number of protons but different neutron counts, leading to different mass numbers.

Writing Atomic Symbols

  • General formula:
    • Mass number (p + n) over atomic number (p) followed by the element symbol.
    • Example for Helium: ( ^4_2He ) (2 protons, 2 neutrons).
  • Example for Uranium:
    • Uranium has 92 protons and 143 neutrons:
    • ( ^{235}_{92}U )

Average Atomic Mass

  • Definition:
    • The weighted average of the isotopes of an element.
  • Example Calculation:
    • For Carbon: Isotopes ( ^{12}C (98.93\%) ) and ( ^{13}C (1.07\%) ):
    • ( ext{Average mass} = (12.0000 ext{ u})(0.9893) + (13.0034 ext{ u})(0.0107) = 12.01 ext{ u} )

Historical Models of the Atom

  • Dalton’s Atomic Theory (1808): Focused on indivisible atoms.
  • Plum Pudding Model (1904): Proposed electrons scattered within a positive matrix.
  • Bohr Model (1913): Suggested electrons are in orbits around the nucleus.
  • Quantum Model (1920s): Explains electrons with properties of particles and waves.

Introduction to Ions

  • Definition: An ion is an atom or group of atoms with an overall charge due to loss or gain of electrons.
  • Positive Ion (Cation): Formed when an atom loses electrons.
    • Example: Lithium ion has 3 protons and 2 electrons (net charge of +1).
  • Negative Ion (Anion): Formed when an atom gains electrons.
    • Example: Fluoride ion has 9 protons and 10 electrons (net charge of -1).
  • Sulfide Ion Example:
    • Sulfur with atomic number 16:
    • Sulfide ion (charge -2) contains 16 protons and 18 electrons.

Summary of Atomic Concepts

  • Atomic identity is determined by protons.
  • Mass number is the total of protons and neutrons (p+n).
  • Isotopes have the same protons but different neutrons.
  • The periodic table is organized by atomic number and average atomic mass.

These notes encapsulate comprehensive information concerning atomic theory, conservation laws, elemental properties, and the structure of atoms as discussed within this chapter. This guide is intended as a detailed resource for further studies in introductory chemistry.