UNIT 8: LEWIS STRUCTURES

Bonding/LE

Conceptual stuff

Lewis Structures

Bond Polarity

More potential energy means more unstable... more wanting to tend towards stability.

More exothermic means harder to reverse! Think about flipping the sign.

Coulomb's law is always being considered!!!

Ionic bonding: effectively transferred electrons. B/n metals and nonmetals (electroneg. difference of >1.7-2)

Covalent Bonding: Polar to nonpolar. More akin to sharing electrons. (electroneg. diff: 0-1.7)

Metallic bonding: How metals bond. Like a sheet of cations and a sea of electrons. Conductive, shiny, etc.

To draw a lewis structure:

  1. Add up all the valence electrons... adjusting w/charge accordingly.
  1. Position w/central atom (CH4 has C as central, HCN has C as central.) Draw preliminary single bonds to other atoms.
  1. Making sure to subtract electrons for bonds you just made, fill outer atom's octets then inner as much as you can (even if it "overflows" it so to speak.)
  1. If it truly overflows or is impossible to reach octets for the central atom if it would violate outer atom's octets, just leave it (Exceptions.) Otherwise, add double/triple bonds until octets are full.

FORMAL CHARGE:

  1. All unbonded, as well as in a bond but closest to the atom are assigned to that atom. (basically electrons "surrounding")
  1. Valence - assigned = formal charge.

Pick the structure where formal charges are closest to zero, or negative formal charges reside on more electroneg. atoms (surrounding atoms.)

RESONANCE STRUCTURES

Basically, a structure where octets can be fulfilled in multiple ways... and there isn't one "most important" structure.

The placement of the atoms in alternative but completely equivalent Lewis structures is the same, but the placement of the electrons is different

Use double arrows to express them (b/n resonances)

Bonds are actually blended average of all structures.

Exceptions (to octet)

  1. Odd # of electrons- rare.
  1. More than 8 on central atom- only happens w/atoms at least at period 3, because they have d orbitals they can fill
  1. Less than 8- happens w/ Be and B as central. Technical rule: if formal charge is negative on central or pos on outside, something is wrong; Boron/Be actually more want to lose electrons

More bonds b/n atoms (double, triple) the stronger/shorter the bond.

ENERGY ALWAYS RELEASED WHEN BONDS ARE FORMED, AND ALWAYS ABSORBED WHEN BONDS ARE BROKEN

Bonds are a on a spectrum from nonpolar to more polar (totally polar is ionic.) From more similar to sharing, to effectively transferred.

When the differences in electronegativity, or tendency to attract electrons in a bond, is high enough, it is more like a transfer and more polar (THINK.)

LATTICE ENERGY: Energy required to break apart a mole of ionic solid (lattice.) Even though ionization energy is usually more endothermic than electron affinity is exo, the attraction of opposite charges allows lattice to be favorable.

When thinking LE/bond favorability, always coulomb's law; also low ionization energy and high electron aff. is favor. bond.

CHARGE matters more than radii in coulomb's law; radii changes negligibly compared to charges.

Electronegativity increases with electron affinity and ionization energy.

Bond enthalpy

Bond enthalpy is the energy required to break a bond (endothermic.)

Nuclei's attraction to other electrons must outdo its repulsion from each other

Therefore, total enthalpy of a reaction is (bond enthalpies of bonds in reactants broken) - (bond enthalpies of bonds in products formed.)

Whenever oxidation state of metal increases, covalent character of bond increases

Exothermic is tending towards stability.

Lewis structures --> formal charge --> if there are multiple w/equivalent formal charge, resonance