Chemical Patterns - Summary Notes

The Periodic Table

  • Elements and compounds:
    • Made of atoms.
    • Arranged in the periodic table by increasing atomic number (usually coincides with increasing atomic mass).
  • Periods:
    • Rows in the periodic table.
    • The period number indicates the number of energy levels/shells.
  • Groups:
    • Columns in the periodic table.
    • Elements in the same group have similar properties and the same number of electrons in their outermost shell.
      • Group 1: Alkali metals
      • Group 2: Alkaline Earth metals
      • Group 17: Halogens
      • Group 18: Noble gases (unreactive due to full outer shells)
  • Metals:
    • Located on the left and middle of the periodic table.
  • Non-metals:
    • Located on the right of the periodic table.
  • Transition metals:
    • Located in the middle of the periodic table.

Properties of Metals and Non-metals

  • Metals:
    • Lustrous, conduct heat and electricity, hard (except Hg), high density, high melting and boiling points (except Hg), malleable, ductile, sonorous.
    • Metallic bonding:
      • Atoms lose electrons to become positively charged ions (cations) in a 3D lattice.
      • Electrons are delocalized and move freely through the structure (sea of delocalized electrons).
  • Non-metals:
    • Dull or glassy, brittle, poor conductors of electricity and heat, low melting points.
    • Many are gases at room temperature.

Representing Atoms & Isotopes

  • Subatomic particles:
    • Protons: positive charge, in the nucleus, relative mass = 1, equal to the atomic number.
    • Neutrons: neutral charge, in the nucleus, relative mass = 1, calculated by: atomic mass - atomic number.
    • Electrons: negative charge, in energy levels/shells, relative mass = 0, equal to the atomic number.
  • Atomic Symbol:
    • One or two letters representing an element.
    • First letter is capitalized, second is lowercase.
  • Relative Atomic Mass:
    • Average mass of an element relative to Carbon-12, no units.
    • Takes into account naturally occurring isotopes.
  • Isotopes:
    • Atoms with the same number of protons but different numbers of neutrons.
    • Same atomic number, different mass numbers.
  • Isotopic Symbols:
    • Mass number (protons + neutrons) at the top, atomic number (protons) at the bottom.

Electron Structure

  • Electronic structure/configuration:
    • Electrons orbit the nucleus in specific energy levels called shells.
      • 1st shell: up to 2 electrons.
      • 2nd shell: up to 8 electrons.
      • 3rd shell: up to 8 electrons.
    • The number of electrons an element has dictates its reactions and uses.
    • Elements in the same group have the same number of electrons in their outer shell and react similarly.
    • Elements in the same period have electrons occupying the same number of shells.
  • Noble Gases:
    • Group 18, have full outer shells of electrons and are unreactive.

Reactivity of Metals

  • Reactivity trends of the Periodic Table:
    • Atoms with full outer shells are stable.
    • Atoms with one electron in their outer shell are very reactive (easily lose the electron).
    • Atoms with seven electrons in their outer shell are also very reactive (easily gain one electron).
  • Reactivity trends when elements lose an electron (cations):
    • Down Group 1, atoms have additional electron shells.
    • Outer electrons are further from the nucleus, decreasing electrostatic attraction, lost easily.
    • Group 1 and 2 metals get more reactive down the group.
  • Reactivity trends when elements gain an electron (anions):
    • Down Group 17, atoms have additional electron shells.
    • Outer electrons are further from the nucleus, attracting electrons less easily.
    • Group 17 halides get less reactive down the group.
  • Metals and Metal Salts:
    • A + BC → AC + B, where A and B are metals. Metal A displaces metal B if it is more reactive

Ions & Ionic Bonding

  • Formation of Ions:
    • Atoms lose or gain electrons to achieve a full outer shell.
    • Losing electrons results in a positive charge (cations).
    • Gaining electrons results in a negative charge (anions).
  • Ions & Charges:
    • Compare the number of protons to electrons present for that atom.
    • If an atom loses 2 electrons, it will form an ion with a charge of +2.
    • If an atom gains 2 electrons, it will form an ion with a charge of -2.
  • Excited Electrons:
    • Electrons jump to other shells when given energy.
    • They release energy in the form of light when they return to the ground state.
  • Ionic Bonding:
    • Electrons are transferred.
    • Atoms losing electrons become positively charged.
    • Atoms gaining electrons become negatively charged.
    • Oppositely charged atoms attract each other (electrostatic attraction).
    • Usually between metals and non-metals forming lattice structures.
  • Ionic Formula:
    • Drawing Method
    • “Balance the Charge Method”
    • “Swap & Drop Method”

Polyatomic Ions

  • Simply treat the entire group of atoms as though it were a single atom

Covalent Bonds & Molecules

  • Covalent compounds are formed when non-metal atoms react together.
  • Non-metals share outer shell electrons to achieve a full outer shell.
  • A pair of shared electrons is a covalent bond.
  • Can be single, double, or triple bonds.
  • Representing Covalent Molecules:
    • Full Bonding Diagrams
    • Outer Shell Bonding Diagrams
    • Lines

Naming Covalent Molecules

  • Name the non-metal furthest to the left on the periodic table by its elemental name.
  • Name the other non-metal by its elemental name but with an -ide ending.
  • Use the prefixes mono-, di-, tri-…. to show how many of that element is in the molecule.