Chemical Bonding - Chemistry

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Chemical Bonding

  • 3 types:

    • Ionic Bonding

    • Metallic Bonding

    • Covalent Bonding

  • cation (+)

    • pronounced cat-ion

  • anion (-)

    • pronounced an-ion

Ionic Bonding

  • attraction that forms between a (+) ion and a (-) ion

  • results from a transfer of e-


Ionic Compounds

  • held together by an ionic bond

  • exists in a 3D “crystal lattice” (repeating pattern)

  • ratio of ions is the chemical formula

  • common properties:

    • very brittle

      • aligning (+) and (+) or (-) and (-) causes a split/ break

    • very high melting point

      • melting requires break ionic bonds to liquify the lattice

    • form electrolyte solutions when dissolved

      • separated ions can conduct electricity

Metallic Bonding

  • attraction between a metal atoms nucleus (+) and another atoms delocalized valence e-

    • the e- orbit all the atoms

    • “sea of electrons” hold all the atoms together


Metal Properties

  • malleable (change shape without it breaking)

  • excellent conductor of electricity

    • valence e- are mobile and can easily move

  • excellent conductor of heat

    • easy for atomic vibrations to travel from atom to atom

Covalent Bond

  • attraction between the nucleus of one atom and a valence e- from another atom

  • results in a sharing of e0 between non metals

  • valence e- orbits 2 atoms

Drawing a Molecular Structure

  1. add up the total valence e- on the molecule

  2. put the least electronegative atom in the center and connect the other atoms by a single bond

  3. add remaining e- as “lone pairs”

  4. rearrange long pairs and bonds to make sure each atom has the octet (only if necessary)

Molecular structure example

Valence Shell Electron Pair Repulsion (VSEPR) theory

  • Electron Domain Geometry

    • 2 bonds = linear

    • 3 bonds = trigonal planer

    • 4 bonds = tetrahedral

  • Molecular Geometry

    • linear:

      • 0 lone pairs = linear

    • trigonal planer:

      • 0 lone pairs = trigonal planer

      • 1 lone pair = bent

    • tetrahedral:

      • 0 lone pairs = tetrahedral

      • 1 lone pair = trigonal pyramidal

      • 2 lone pairs = bent

Steric number = number of bonds

Octet exceptions

  • Electron Domain Geometry

    • 5 bonds = trigonal bipyramidal

    • 6 bonds = octahedral

  • Molecular Geometry

    • trigonal bipyramidal:

      • 0 lone pairs = trigonal bipyramidal

      • 1 lone pair = seesaw

      • 2 lone pairs = T-shaped

      • 3 lone pairs = linear

    • octahedral:

      • 0 lone pairs = octahedral

      • 1 lone pair = square pyramidal

      • 2 lone pairs = square planer

      • 3 lone pairs = T-shaped

      • 4 lone pairs = linear

Resonance

  • when e- spend time between more than 2 atoms

Polar Covalent Bond

  • bond in which electrons pairs are shared unequally

  • due to atoms in the bond having different electronegativity

  • the bond has partially (-) and (+) ends

    • partial = δ (lowercase delta)

Electronegativity difference

difference ≤ .4 → nonpolar covalent bond

.4 < difference ≤ 1.8 → polar covalent bond

difference > 1.8 → ionic bond

  • ONLY IF THE ATOMS ARE POLAR

    • δ - = more electronegative atom

    • δ + = less electronegative atom


Determining if a Molecule is a Dipole

Dipole = polar molecule

  1. Must have at least 1 polar bond

  2. Molecule must be asymmetrical (if there are lone pairs on the central atom, it’s probably asymmetrical)