Atoms and the Periodic Table Study Notes

Chapter 2: Atoms and the Periodic Table

Review

  • Behavior of electric charges: attraction and repulsion.

Learning Objectives

  • Explain the major assumptions of atomic theory.
  • Identify subatomic particles and their properties.
  • Write the symbols for different isotopes of an element using mass numbers and atomic numbers.
  • Classify an element as a metal, nonmetal, or metalloid.
  • Describe chemical behavior based on group membership.

Atomic Theory

  • Atom: The smallest unit of an element that retains the properties of that element.
  • Dalton’s Atomic Theory:
    1. All matter is made up of atoms.
    2. All atoms of an element are the same, and they are different from atoms of other elements.
    3. Atoms of multiple elements combine to form compounds; a particular compound is always made up of the same types of atoms in the same proportions.
    4. Chemical reactions only involve the rearrangement of atoms, not their destruction or creation.

Atomic Structure

  • Subatomic Particles:
    • Neutron (n, n0): Mass ~1 amu.
    • Proton (p, p+): Mass ~1 amu.
    • Electron (e–): Mass ~0.0006 amu.

Atomic Numbers

  • Atomic Number: The number of protons in atoms of an element. It is unique to each element (e.g., carbon always has 6 protons).
  • The atomic number also equals the number of electrons in a neutral atom.
  • Mass Number: Number of protons plus number of neutrons.
  • Isotopes: Atoms with the same number of protons but different numbers of neutrons.
    • Example: “carbon-12” or “C-12” has an atomic number of 6 and a mass number of 12.

Atomic Numbers Practice

  • Provide missing information: Element name/symbol, atomic number, mass number, and number of protons/neutrons/electrons.
  • Example: lithium-6, atomic # 3; mass # 6; protons 3; neutrons 3; electrons 3.
  • Other examples: Chlorine-37, an atom with 10 protons and 10 neutrons, oxygen with 10 neutrons, an atom with 19 protons and a mass number of 31 (incomplete info).

Organizing the Elements: The Periodic Table

  • The periodic table organizes elements by various properties:
    • Metals: shiny, ductile, good conductors, high melting points, and high densities.
    • Nonmetals: poor conductors, low melting points, and low densities.
    • Metalloids: semiconductors; possess properties between metals and nonmetals.
    • Groups:
    • 1: Alkali metals (highly reactive with water).
    • 2: Alkaline earth metals (less reactive than alkali metals).
    • 13-16: Various groups including Pnictogens, Chalcogens and Halogens.
    • 17: Halogens (highly reactive).
    • 18: Noble gases (unreactive).

Organization of Elements by Physical Properties

  • Elements are sorted into groups based on similar chemical behaviors and properties.

Groups in the Periodic Table

  • Example elements with their respective atomic number and atomic weight:
    • Hydrogen (H): Atomic Weight = 1.008, Atomic Number = 1.
    • Lithium (Li): Atomic Weight = 6.94, Atomic Number = 3.
    • Beryllium (Be): Atomic Weight = 9.0122, Atomic Number = 4.
    • Sodium (Na): Atomic Weight = 22.990, Atomic Number = 11.
    • Magnesium (Mg): Atomic Weight = 24.305, Atomic Number = 12.

Behavior of Elements Based on Group Membership

  • Elements are often grouped based on their reactivity or common properties. For example:
    • Alkali metals react strongly with water.
    • Halogens are highly reactive nonmetals.
    • Noble gases are typically unreactive.

Electron Energies

  • In general, objects can possess any value of energy (like standing at any height on a ramp).
  • In contrast, electrons in an atom are quantized, having only certain allowed energy values (like standing on specific steps of a staircase).

Organization of Electrons in an Atom

  • Electrons reside in orbitals. Each orbital can hold 2 electrons with opposite spin (spin can be either up or down).
  • Orbitals are organized into subshells:
    • s: 1 spherical orbital.
    • p: 3 dumbbell-shaped orbitals.
    • d: 5 orbitals.
    • f: 7 orbitals.
    • Increasing energy levels can accommodate electrons in shells.

Electron Configuration of an Atom

  • Electrons fill orbitals in order from lowest to highest energy.
  • Configuration is indicated by the address of each electron and the number of electrons at that address.
  • Noble gas configuration (e.g., [He]2s²) is a shorthand notation for long electron configurations.

Valence Electrons

  • Valence Electrons: Electrons located in the outermost shell (highest principal quantum number, n).
  • They primarily determine the chemical reactivity and properties of elements.
  • Example configurations:
    • Lithium (Li) has 1 valence electron: [He] 2s¹.
    • Sodium (Na) has 1 valence electron: [Ne] 3s¹.
    • Potassium (K) has 1 valence electron: [Ar] 4s¹.

Lewis Symbols

  • Lewis Symbols (or electron-dot symbols) are used to represent an atom and its valence electrons. Dots represent electrons, which should be arranged such that they are spread out before any pairing occurs.

Practice Tasks

  • Work in groups to identify the group number of the listed elements, their group type (metal/nonmetal/metalloid), and predict their properties.
  • Conduct electron configuration exercises to write out configurations for specific elements (e.g., Magnesium, Aluminum).

Reminders

  • Homework and assignment due dates for following chapters and reports:
    • Chapter 1 Homework due 2/4 at 11:59 PM
    • Chapter 2 Homework due 2/11 at 11:59 PM
    • Experiment 1 Report due next lab period.