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5.111 Lecture Summary #10 Monday, September 29, 2014
Readings for today: Section 2.5 – 2.8 Lewis Structures (Same sections in 5th^ and 4th^ ed.) Read for Lecture #11: Section 2.9 – Radicals and Biradicals, Section 2.10 - Expanded Valence Shells, Section 2.11 - Group 13/III Compounds (Same sections in 5th^ and 4th^ ed.)
Topics: I. Lewis structures II. Formal charge III. Resonance structures
I. LEWIS STRUCTURES G.N. Lewis (American scientist, 1875-1946). Twenty years prior to the development of quantum mechanics, Lewis recognized an organizing principle in bonding. Namely that:
The key to covalent bonding is electron sharing , such that each atom achieves a valance shell (noble gas configuration).
OCTET RULE: electrons are distributed in such a way that each element is surrounded by eight electrons, an octet. Each dot in a Lewis structure represents a e-.
EXCEPTION WITH H: special stability is achieved with electrons.
Each valence e-^ in a molecule can be described as a bonding or a lone-pair electron. For Cl in HCl
- bonding electrons
- lone-pair electrons or lone pairs
Lewis structures correctly predict electron configurations 90% of the time. Our other option: solve the Schrödinger equation.
PROCEDURE FOR DRAWING LEWIS STRUCTURES
- Draw a skeleton structure. H and F are always terminal atoms. The element with the lowest ionization energy goes in the middle (with some exceptions).
- Count the total number of valence electrons. If there is a negative ion, add the absolute value of total charge to the count of valence electrons; if positive ion, subtract.
- Count the total # of e-s needed for each atom to have a full valence shell.
- Subtract the number in step 2 (valence electrons) from the number in step 3 (total electrons for full shells). The result is the number of bonding electrons.
- Assign 2 bonding electrons to each bond.
- If bonding electrons remain, make some double or triple bonds. In general, double bonds form only between C, N, O, and S. Triple bonds are usually restricted to C, N, and O.
- If valence electrons remain, assign them as lone pairs, giving octets to all atoms except hydrogen.
- Determine the formal charge.
EXAMPLES
Hydrogen cyanide (HCN) Cyanide ion (CN-)
- skeletal structure. (Atom in the middle for HCN is )
of valence e-s. (Don’t forget charges)
of e-s for each atom to have a full valence shell.
of bonding e-s.
- Assign 2 bonding electrons per bond.
- remaining bonding electrons?
- remaining lone pair e-s?
- determine formal charge (we will come back to this in a minute).
II. FORMAL CHARGE (FC)
Formal charge is a measure of the extent to which an atom has gained or lost an
in the process of forming a covalent bond.
If two valid Lewis structures have the same absolute value of formal charges, the more
stable structure is the one with a negative formal charge on the more
atom.
CH 3 NHO-^ For “chain” molecules, atoms usually written in order. Terminal atoms usually follow
CH 3 is a group.
These groups are always
terminal. (^) the atom to which they are attached.
H -1^ H^ - H C N O^ H^ C^ N^ O^ H H H^ H
Zero FC on all other atoms Zero FC on all other atoms χ: F > O > N > C energy structure
III. RESONANCE STRUCTURES
For certain molecules, more than one Lewis structure is needed to correctly describe the valence electron structure of the molecule.
For example, consider the Lewis structure(s!) of ozone, O 3.
structure 1 structure 2
skeletal structure
valence e–s: 3(6) =
full shell e–s: 3(8) =
bonding e–s: – =
assign bonding e –^ s
remaining bonding e –^ s: 2
remaining valence e–^ s (assigned as lone pairs): 12
formal charges:
Structure 1 Structure 2 FCO A = FCO A = FCO B = (^) FCO B =
FCO C = FCO C =
We might expect one short O=O bond and one long O-O bond, but experimental evidence demonstrates that the two bonds are.
Thus, the two structures are equivalent. A better model is to blend the structures as
denoted with the brackets and arrows below, a resonance hybrid.
Electrons in resonance structures are.
Electron pairs are shared over several atoms, not just two.
Resonance structures are two (or more) structures with the same arrangement of
ATOMS, but a different arrangement of ELECTRONS.
