Chemical Bonding

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Chemical Bonding

• Chapter 7

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Chemical bonding

• ionic bond an electrostatic attraction between
  ions of opposite charge
• covalent bond sharing electrons between atoms
• metallic bond bonding occurs throughout the
  substance, electrons flow freely throughout the
  metal
• (metals make good conductors)

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Ionic bonding

• ions group together in a lattice formation.
  There is no single molecule of an ionic
  compound.
• What kind of energy changes are involved in
  formation of an ionic compound?
• Na (g) Cl (g) ?? NaCl (g)

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What trends are there in lattice energies?

• the potential energy between two particles.
• Qcharge on particle 1 or 2, ddistance between
  the particles, kconstant.
• So, if magnitude of Q1 and Q2 increases, the
  energy _______.
• If d increases, the energy ______, but this
  change is not as much as when the charges are
  changed.

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Crystal lattice energies

• Determine the order of the crystal lattice
  energies for the following ionic compounds, from
  lowest to highest
• MgO, KBr, ScN
• KBr (671 kJ/mol) lt MgO (3795 kJ/mol) lt ScN (7547
  kJ/mol)

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Lewis symbols

• a pictoral representation of the valence
  electrons of an atom
• valence electrons vs. core electrons
• examples

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Ionic compounds

• In binary ionic compounds (i.e., no polyatomic
  ions involved), the atom to be the anion is
  formed by swiping the electrons of the atom to
  be the cation.
• Example NaCl

More examples
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Covalent bonds

• Instead of losing or gaining electrons, atoms
  share electrons so that a bond is formed
  between them.

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Lewis formulas in covalent bonding

• Lewis formulas show how atoms share electrons
  between each other to form a bond.
• Examples H2O, CH4, NH3

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Octet rule

• Atoms tend to gain, lose, or share electrons
  until they are surrounded by EIGHT valence
  electrons.
• Why eight?
• The octet rule is not a law, and there are
  several exceptions we will discuss later.

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Rules for drawing Lewis structures

• Sum all valence electrons from all atoms
• Arrange the atoms
• linear molecules (normally two or three), formula
  sometimes shows atoms from right to left
• central-grouped atoms center atom normally
  written first.
• Least electronegative element in the center
• complete octets around all atoms bonded to
  central atom
• remember H has only 2 electrons
• place all leftover electrons on central atom
  (even if more than an octet results)
• try multiple bonds if the central atom doesnt
  have an octet using single bonds
• Examples H2O, CO, NH3

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Formal charge

• A bookkeeping tool used to help determine the
  best structure between two or more structures
  that follow the octet rule. (NOT the same as
  oxidation numbers)
• valence electrons
• assigned electrons all the unshared electrons
  (nonbonding electrons) ½ all the bonding
  electrons
• FC Valence electrons assigned electrons
• Sum of FC of all atoms must charge.
• Rule The structure resulting in lowest formal
  charge on each atom wins.
• Example structures of N2O

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Resonance structures

• structures meeting both octet rule and formal
  charge requirements.
• Resonant structures are equivalent, but they are
  also equally wrong.
• The actual structure is in between all
  resonance structures.
• Examples O3, NO3-, SO3

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Exceptions to the octet rule

• Odd number of electrons in molecule.
• Molecules where an atom has less than an octet.
• Molecules where atom has more than an octet.
• Exception 1 examples NO, ClO2
• Exception 2 examples BeCl2, BF3

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Exception 3 examples (most common)

• PCl5, SF6, PO43-
• Central atom expands into its d shell orbitals.
• For this to occur, central atom must be 3rd row
  or higher.

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Polar and non-polar bonds

• Not all atoms share electrons equally in the
  atom.
• Electronegativity is an indicator of how much one
  atom will dominate the possession of electrons.
• What is electronegativity?
• The difference in electronegativities (?EN) lets
  us know how polar a molecule is

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Dipole moments

• Dipole moments occur in polar molecules.
• The larger the ?EN, the larger the dipole moment.
• The dipole moment points toward the more
  electronegative element.
• Molecules with more than two atoms are more
  complicated (well talk about that in Chapter 8).
• Examples