CHEM 1040 Lecture Notes - Lecture 2: Lewis Structure, Molecular Geometry, Ionic Bonding
16 May 2020
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1
– M olecular Structure –
LEWIS STRUCTURES (Refer to page 330 and Section 9.6)
Notes & Hints:
1. Chemical properties mostly determined by number of electrons in
outermost shell of atom (i.e., ______________________).
2. Lack of reactivity of inert gases is associated with filled outer s and p shells.
3. Lewis Convention: chemical symbol stands for nucleus and inner electrons;
outermost electrons are represented by dots (or a line signifies an electron
pair), e.g., Na⋅ C | F⋅
4. Lewis diagrams only useful when only have s and p electrons
(not great for d electrons).
5. The first octet period is
6. We now know the Lewis diagrams for ALL atoms in the eight main groups
within the periodic table!
7. To write a Lewis diagram, need to know either electronic configuration or
position in periodic table.
8. Atoms tend to achieve a complete outer shell (octet status) either by:
a) losing or gaining e–’s (ionic bonding) e.g., Na → Na+, Cl → Cl
–
b) sharing (covalent bonding).
Simple rules for drawing covalent Lewis Diagrams
(refer to page 7–4 in your lab manual)
1. Determine the total number of valence electrons, V.
2. Draw skeleton structure using single bonds.
3. Subtract number of electrons used thus far from V and then distribute these
extra electrons so each atom achieves octet status (access to 8 electrons).
[Note: H can only get 2 electrons and F only forms single bonds.]
4. If too few electrons to satisfy all, use double or triple bonds until each [non-
H] atom has 8 electrons.
5. If too many electrons , add extras to central atom as lone pairs (expanded
octet allowed when central atom has n ≥ 3 (row 3 or greater).
Note: We are just "electron accounting" at present and NOT implying geometry
(that comes later).
NOTES:
2
EXAMPLES:
NF3
C2H4
N2O (N-N-O)
CO32–
CH3NH3+
FORMAL CHARGES: (refer to page 7–6 in your lab manual)
To deduce the formal charge on an atom:
a. Assign 1 electron from each SHARED electron
pair (i.e., one per single
bond, two for double, etc.) and then ADD to this the total number of
UNSHARED pairs of electrons.
b. If atom has ____________ outer electrons than it would as a neutral atom,
it has a POSITIVE charge.
c. If atom has MORE outer electrons than it would as a neutral atom, it has a
______________ charge.
The SUM of all the formal charges from all the atoms:
i) within a MOLECULE must be ZERO
ii) within an ION must be the _______________ on the ion
NOTES:
3
EXAMPLES:
NH4+
ClO –
(CH3)3NO
RESONANCE:
Sometimes it’s not possible to write a single Lewis diagram which correctly
reflects the known chemistry of the molecule, e.g., SO2
S S
?
O O O O
It’s important to recognize what we mean and what we
DO NOT mean by RESONANCE:
Ø true state of bonding is __________ correctly represented by ANY ONE
Lewis diagram
Ø instead it exists somewhere ______________________ these extremes!
Thus the reason we link the resonance structures by
a double headed arrow.
NOTES:
Examples of Resonance
NO3–
CO32–
SCN–
N3–
