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