bondingpractice.notebook
Lewis Dot Structure
Procedure for Neutral Molecules (CO2)
September 16, 2014
Procedure for Negatively Charged Ions (CO32­)
1. Use the same procedure as before, then as a last step add one electron per negative charge to fill octets. 1. Decide how many valence (outer shell) electrons each atom in the molecule has.
2. If there is more than one atom type in the molecule, put the most metallic or least electronegative atom in the center. Recall that electronegativity decreases as atom moves further away from fluorine on the periodic chart.
3. Arrange the electrons so that each atom contributes one electron to a single bond between each atom.
4. Count the electrons around each atom: are the octets complete? If so, your Lewis dot structure is complete.
5. If the octets are incomplete, and more electrons remain to be shared, move one electron per bond per atom to make another bond. 6. Repeat steps 4 and 5 as needed until all octets are full.
7. Redraw the dots so that electrons on any given atom are in pairs wherever possible.
Sep 4­8:24 AM
Sep 4­8:29 AM
Procedure for Positively Charged Ions (NH4+)
Violations of the Octet Rule
Use the same procedure as outlined above, then remove one electron per postive charge as needed to
avoid expanded octets. When using this procedure for positively charged ions, it may be necessary to have some atoms with expanded octets (nitrogen in this example). For each unit of positive charge on the ion remove on electron from these exanded octets. If done correctly, your final structure should have no first or second period elements with expanded octets.
Odd­electron molecules represent the first violation to the octet rule. Although they are few, some stable compounds have an odd number of electrons in their valence shells. With an odd number of electrons, at least one atom in the molecule will have to violate the octet rule. Examples of stable odd­electron molecules are NO, NO2, and ClO2. Although the O atom has an octet of electrons, the N atom has only seven electrons in its valence shell. Although NO is a stable compound, it is very chemically reactive, as are most other odd­electron compounds.
Sep 4­8:29 AM
Electron­deficient molecules represent the second violation to the octet rule. These stable compounds have less than eight electrons around an atom in the molecule. The most common examples are the covalent compounds of beryllium and boron. Sep 4­8:25 AM
The third violation to the octet rule is found in those compounds with more than eight electrons assigned to their valence shell. These are called expanded valence shell molecules. Such compounds are formed only by central atoms in the third row of the periodic table or beyond that have empty d orbitals in their valence shells that can participate in covalent bonding. For example, beryllium can form two covalent bonds, resulting in only four electrons in its valence shell
Boron commonly makes only three covalent bonds, resulting in only six valence electrons around the B atom. One such compound is PF5. A well­known example is BF3:
Sep 4­8:38 AM
Sep 4­8:39 AM
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bondingpractice.notebook
September 16, 2014
Molecular Geometry
Practice Doing Lewis Dot Structures
a) BeF2
f) BH2–
b) BCl3
g) NI3
c) CCl4
h) ClF4+
d) PBr5
i) SF5–
e) SI6
Sep 4­8:25 AM
Sep 4­9:05 AM
Bonding Angles
Number of regions
Arrangement of
of high electron
regions of high
density around
electron density
central atom
in space
4
tetrahedral
Predicted
bond
Geometry of
Example
molecule
angles
109.5°
CH4, methane
tetrahedral
NH3, ammonia
pyramidal
H2O, water
bent
Example
Predict all bond angles in the following molecules.
a. CH3Cl b. CH3CNl c. CH3COOH
3
2
trigonal planar
linear
120°
180°
H2CO, formaldehyde trigonal planar
C2H4, ethylene
planar
SO2, sulfur dioxide
bent
CO2, carbon dioxide
linear
C2H2, acetylene
linear
Sep 10­1:52 PM
Multiple Bonds
Sep 10­1:54 PM
HCN
CH2O
NHCHCH (hint: cyclo)
NC2H3
Sep 10­1:52 PM
Sep 10­1:52 PM
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bondingpractice.notebook
Formal Charge
September 16, 2014
hybridization
coolness
formal charge = valence electrons ­ electrons assigned to atom in a bond
CH4
WHY?
Why use it? (What does it tell us)
Formal Charge = Valence electrons ­(nonbonding electrons ­ bonded electrons/2)
NH3
ONO
NO2
SF6
H 2S
C 2H 4
Sep 10­1:52 PM
Sep 10­2:11 PM
Sep 16­8:04 AM
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