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