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
1. 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).
2. 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.
3. Count the total # of e-s needed for each atom to have a full valence shell.
4. 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.
5. Assign 2 bonding electrons to each bond.
, 6. 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.
7. If valence electrons remain, assign them as lone pairs, giving octets to all atoms
except hydrogen.
8. Determine the formal charge.
EXAMPLES
Hydrogen cyanide (HCN) Cyanide ion (CN-)
1. skeletal structure.
(Atom in the middle for HCN is )
2. # of valence e-s. (Don’t forget charges)
3. # of e-s for each atom to have a full valence shell.
4. # of bonding e-s.
5. Assign 2 bonding electrons per bond.
6. remaining bonding electrons?
7. remaining lone pair e-s?
8. 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.
2
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
1. 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).
2. 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.
3. Count the total # of e-s needed for each atom to have a full valence shell.
4. 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.
5. Assign 2 bonding electrons to each bond.
, 6. 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.
7. If valence electrons remain, assign them as lone pairs, giving octets to all atoms
except hydrogen.
8. Determine the formal charge.
EXAMPLES
Hydrogen cyanide (HCN) Cyanide ion (CN-)
1. skeletal structure.
(Atom in the middle for HCN is )
2. # of valence e-s. (Don’t forget charges)
3. # of e-s for each atom to have a full valence shell.
4. # of bonding e-s.
5. Assign 2 bonding electrons per bond.
6. remaining bonding electrons?
7. remaining lone pair e-s?
8. 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.
2
