Unit 6 summary
By Nazanin & Brittany H.
Bonding & Molecular Structure
1. The structure of a compound, either ionic or covalent, is the one with having the
lowest potential energy_ that is the one with the greatest thermodynamics stability.
2. Predicting Lewis structure:
Example: What is the Lewis structure for NH2Cl?
a. Put a single bond :
HN Cl

H
b. Calculations :
Total valence e : 5+1+1+7= 14
e used by single bonds:
6
Have(remaining) :
8
Need to complete the octet: 8
c. If need = have, then add lone pairs.
If need > have, then add bonds # of multiple bonds = (need – have ) /2
If need < have, then add extra lone pairs on the central atom.
(Refer to P387, Table 9.4)
3. Exceptions to Octet rule:
a. Incomplete Octets:
Some atoms can have a complete outer shell with less than eight electrons. For example, hydrogen
can have a maximum of two electrons, and beryllium can be stable with only four valence
electrons, as in BeH2:
H  Be  H
b. Expanded Octets:
In molecules that have d subshells available, the central atom can have more than eight valence
electrons. Examples: PCl5, SF4, and SF6
c. Odd numbers of electrons:
Molecules almost always have an even number of electrons, allowing electrons to be paired, but
there are some exceptions, usually involving nitrogen. For example, NO and NO2.
NO:
NO
4. Resonance structures are a way to represent bonding in a molecule or ion when a single
Lewis structure fails to describe accurately the actual element structure.
Example:
[ ONO ]
[ ONO ]
5. Formal charge of an atom = (# of valence e of that atom)  (# e ‘near’ that atom in
the structure). For example the formal charge of N in the above NO2 structure is ( 5  5 ) 0.
1
Given different alternatives for the structure of a molecule, the one which minimizes the formal
charge is preferable. The sum of the formal charges in an structure is the charge of the ion.
6. Diamagnetism= no unpaired electrons, not magnetic at all.
Ferromagnetism = permanent magnets: iron, cobalt, and nickel.
Paramagnetism = result of the unpaired electrons in the orbitals of an atom.
7. Quantum numbers: the position of the electrons in relation to the nucleus.
a. n = the principal quantum number, 1,2,3,… The number of “electron shell.”
b. L = the angular momentum quantum number, 0,1,2,3,…,n1
Value of L
0
1
2
3
Corresponding sub shell label
s
p
d
f
c. mL = the magnetic quantum number that describes the orientation of the orbital in
space.
Subshell
Value of mL
s (L=0)
p (L=1)
d (L=2)
0
-1,0,1
-2,-1,0,1,2
d. ms = the electron spin magnetic quantum number, +1/2, 1/2
8. Bond order = ½[(# bond e)  (# anti-bond e)]
9. Orbital hybridization: new sets of orbitals, called ‘hybrid orbitals,’ could be created by
mixing s, p, and/or d atomic orbitals on an atom.
The number of hybrid orbitals is the same as the number if atomic
orbitals used in their structure.
The hybrid orbitals of an atom is sp3 when the atom has 4  bonds
and/or electron pairs; it is sp2 when it has 3  bonds and/or relectron pairs, and it is sp when it
has 2  bonds and/or electron pairs.
10. Periodic trends :
 Electronegativity increases by moving across a period from left to
right and going up in a family.
Atomic radius increases by moving across a period from right to
left and going down in a family.
Ionic size: cations are smaller than the original neutral atom, and
anions are bigger than the neutral atom.
Ionization energy increases by moving from left to right across a
period and moving up on a group.
11. Polarity:
2
Any bonds
Polar?
No
Molecule is
non-polar
Yes
Molecular is
non-polar
Yes
Distributed
Symmetry?
No
Molecular is
Polar
3