Transcript Bonding & Molecular Structure

Unit 6
Nazanin Ashourian
Brittany Haynes
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. Exeptions 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 excptions, usually involving nitrogen. For example, NO
and NO2.
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.
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.
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, nickel.
Paramagnetism = result of the unpaired electrons in the orbitals of an atom.
7.
Quantum numbers:
electrons in relation to the nucleus.
the position of the
 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. = the magnetic quantum number that describes the orientation of the
orbital in space.
Subshell
Value of mL
s (L=0)
0
p (L=1)
-1,0,1
d (L=2)
-2,-1,0,1,2

b.
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.