DL

Unit 5 - Chemical Bonding

Ionic Bonding

  • Determining if a bond is ionic
    • If difference in electronegativity of two molecules is greater than 1.7, then they form an ionic compound
    • Metals bonding w/ non-metals most likely form ionic bonds
    • Metals tend to form positive ions while non-metals tend to form negative ions
  • Crystals - See unit 4
    • Crystal Lattice - 3d model of arrangement of atoms in a crystal
  • Formula Units - Smallest repeating unit in a compound
  • Polyatomic Ions - A unit made of multiple ions bonded covalently
    • Has an individual charge that is evenly distributed throughout the atom
    • Can combine with other ions to create ionic compounds
    • Treat it as a normal molecule w/ a charge
    • When the polyatomic ion requires a subscript, it is written in parenthesis
    • Mostly anions(negative)
    • Lots differ only in # of oxygen or hydrogen
  • Lattice Energy - Energy released when atoms in the gas phase combine to form crystalline structures
    • Larger for smaller ions/ions with larger charges
    • Can be used to determine the bond strength between atoms
  • Properties of Ionic Compounds
    • High boiling/melting points (solid crystals at room temperature)
    • Hard
    • Somewhat soluable
    • Low conductivity as solids, high conductivity as liquids/gases or when dissolved

Covalent Bonding

  • Typically forms between 2 non-metals
  • The two atoms “share“ electron(s)
    • Electrons that are shared are called bonding electrons, while ones that aren’t shared are called non-bonding/lone pair electrons
  • Sigma Bonds - When the s, p, or d orbital of one atom overlaps with the s, p, or d orbital of another atom
  • Single Bonds - A covalent bond that shares 2 electrons; typically represented as a single line
  • Non-Polar Bonds
    • Forms between two atoms with <=0.4 difference in electronegativity
    • Electrons are shared equally between the two atoms
  • Polar Bonds
    • Forms between two atoms with >=0.5 difference in electronegativity
    • Electrons are not shared equally (they spent more time around one atom than the other)
    • Develops partial charges
  • Lewis Structures
    • Displays the chemical symbol surrounded by the atom’s valence electron(s)
    • Can be used to show bonds between atoms
  • Special Bonds
    • Double Bonds - Two atoms share four electrons
    • 1 Sigma Bond + 1 Pi Bond
    • Triple Bonds - Two atoms share six electrons
    • 1 Sigma Bond + 2 Pi Bonds
    • Second Pi Bond 90 degrees shifted from the first
    • Pi Bonds - Overlapping of P-orbitals, which allows additional electrons to be shared
  • Resonance Structures - When 2+ Lewis structures can be drawn from a chemical formula
  • Defining Properties
    • Don’t ionize in solutions
    • Don’t conduct electricity/heat
    • Low melting/boiling points
    • Gases/liquids at room temperature

Metallic Bonding

  • Properties of Metals
    • Large atoms
    • Low electronegativity
    • Low ionization energy
    • Electrons roam freely → They’re shared among all metallic nuclei → some electrons end up delocalized(electron that doesn’t belong to any individual atom in a crystal)
    • Malleable; delocalized electrons → easy to modify shape
    • Ductile; delocalized electrons → bendable
    • Luster; caused by movement of electrons between bands
  • Molecular Orbitals - A larger orbital made from the combination of 2 or more orbitals
    • Multiple of these combine to form metal “bands“
  • Alloy - Homogenous mixture of multiple metals

Nomenclature of Ionic Compounds

  • Oxidation State - The charge of an element after bonding
    • Transition Metals typically form positive oxidation states
    • Non-metals typically form negative oxidation states
  • Naming Ionic Compounds
    • Name of cation followed by name of anion
    • Metal cations - name of metal
    • Non-metal anions - replace ending with -ide

Nomenclature of Covalent Compounds

  • Name of element further to the left first, second element ends with -ide
  • Prefixes for number of ions
    • Mono, di, tri, tetra, etc
  • No roman numerals
  • Some compounds have common/trivial names that are also used for them
  • Some hydrogen compounds are also exceptions and have names that don’t indicate # of ions
  • Acids - Increases the # of H+ ions in aqueous solution
    • Generally contain H as the only positively bonded ions
    • Turn blue litmus paper red
    • Tastes sour
    • Naming binary acids (acids w/ two atoms)
    • hydro + root of nonmetal + ic + “ “ + acid
    • Naming oxyacids (acids w/ polyatomic atoms that have O in them)
    • -ate → -ic
    • -ite → -ous
  • Bases - Increases the # of OH- ions in aqueous solution
    • Ionic bases have OH- in their chemical formula, while covalent bases don’t
    • Turn red litmus paper blue
    • Taste bitter
    • Feel slippery/soapy
    • Naming bases
    • Covalent are usually amines
    • Based on # of groups attached to the N

Molecular Geometry

  • Valence Shell Electron Pair Repulsion (VSEPR) Theory
    • Pairs of electrons are arranged so that the distance between them is maximized
    • Central atom w/
    • 2 pairs of electrons → linear shape
    • 3 pairs of electrons → trigonal-planar shape
      • + 1 lone pair → trigonal-pyramidal shape
    • 4 pairs of electrons → tetrahedral shape
    • 5 pairs of electrons → trigonal-bipyramidal shape
    • 6 pairs of electrons → octahedral shape
    • Lone pairs also add repulsion energy; causes bent molecular geometry
  • Hybrid Orbitals
    • Orbitals formed by 2+ atomic orbitals combining
    • Formed by valence orbitals mixing together
    • Equal # of hybrid and valence orbitals, but different shape and energies of individual orbitals