Module 5: Bonding, Polarization, & Hybridization

Bond

• Charged electrons attracting nuclei of 2 atoms

Exothermic reaction

• Bond formation releases energy

• Make atoms more stable

Types of bonds

• ionic: formed by sharing electrons bw opposite charges

• covalent: formed when sharing electrons & haver similar electronegativities

Ions types

• ion: charged particle formed by a loss/gain of an electron

• cation: (+) ion attracted to (-) pole/cathode

• anion: (-) ion attracted to (+) pole/anode

Ion charge

• based on amount of electrons gained/loss

Metal ion charge

• lose electrons bc of lower electronegativity

• (+) sign

Non-metal ion charge

• gain electrons bc of higher electronegativity

• (-) sign

Metal + non-metal bonded together is what kind of bond?

• ionic

Ion characteristics

• high boiling/melting point

• does NOT have electric current when SOLID

• conducts electricity when MELTED/DISSOLVED

Metal characteristics

• good heat & electricity conductor

• easily molded

  • plane slide past another w/o bond breaking

• high boiling/melting point

• turns gaseous when melted

• shiny bc absorbs and re-emits light

Ionic characteristics

• poor heat conductor

• conducts electricity when MELTED

• brittle

• lack mobility bc of covalent bonds

Molecular characteristics

• poor heat & electricity conductor

• lacks mobility bc covalent bond

• difficult to mold

Covalent bonds types

• polar: unequal electron sharing — 0.5-1.6 EN

• non-polar: equal electron sharing — < 0.5 EN

• ionic: full electron transfer — > 1.6 EN

VSEPR theory

• electron pairs REPEL

• does not predict # of bonds that determines molecular shape

Polarity determining factors

• molecular shape

• presence/absence of polar bonds

Polar traits

• unequal charged distributed

• µ > 0

Non-polar traits

• no polar bonds

• symmetrical bonds cancel each other

• µ = 0

Dipole moment

• polarity which is based on # value

  • polar: µ > 0

  • non-polar: µ = 0

Hybridization

• mix atomic orbitals to form equal-energy hybrid atoms

• sp₃: tetrahedral

• sp₂: triangular planar

• sp: linear

IntERmolecular forces

• forces holding molecules together in liquid state

  • overcome by increased temp. molecules separate

• weaker than metallic/ionic bonds

Octect rule

• atoms share to reach 8 valence electrons (except H)

Electronegativity calculation

• lower EN - higher EN

Lewis structure

• show valence electron arrangement in molecules

Molecular bond types

• single: 1 shared pair

• double: 2 shared pairs

• triple: 3 shared pairs

Molecular shapes

• linear: 2 groups

• trigonal planar: 3 groups

• tetrahedral: 4 groups

Lone pairs

• non-bonded/unshared electrons

• alters shape of molecule:

  • bent

  • linear

Bond type: BP & strength

• higher BP = stronger

• metallic & ion = highest BP bc ionic bond

• hydrogen (h)-bonding: requires OH, NH, or FH

  • H adds to strength

• dipole-dipole forces: intermediate

• london dispersion forces: weakest

BP: highest to lowest

• metallic

• ionic

• hydrogen bond

• dipole-dipole

• lodon dispersion

Polar substances

• soluble in water

Ionic substances

• soluble in water

Non-polar substances

• not soluble in water

• soluble IF in non-polar solvent

  • like dissolves like

Electron configuration

1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p⁶ 6s² 4f¹⁴ 5d¹⁰ 6p⁶ 7s² 5f¹⁴ 6d¹⁰ 7p⁶