Honors Organic Chemistry
Name: _______________________________
Unit 1 Packet:
Lewis Structures, Resonance, Formal Charge,
VSEPR, Hybridization, and Isomerism
Honors Organic Chemistry
Key Terms For Unit 1
General
Organic Chemistry
Structural Isomerism
Isomerism
Geometric Isomerism
Bonding
Anion
Hund’s Rule
Sigma bond (σ)
Bond Angle
Hybridization
sp hybrid orbital
Bond Length
Ionic Bond
sp2 hybrid orbital
Cation
Line Angle
sp3 hybrid orbital
Covalent Bond
Non-polar covalent bond
Tetrahedron
Double Bond
Octet Rule
Trigonal Planar
Electronegativity
Pauli exclusion Principle
Valence Electrons
Formal Charge
Polar covalent bond
Valence Shell
Full Octet
Pi bond (π)
Valence Shell ElectronPair Repulsion
Hindered Rotation
Resonance hybrid
Nomenclature
Alcohol
Carboxyl group
Functional Group
Aldehyde
Carboxylic Acid
Ketone
Carbonyl group
Ether
Honors Organic Chemistry
Lewis Structures
History:
Purpose:
Process:
CH4
1. Predict the Arrangement of
the atoms:
2. Count up the valence
electrons
3. Connect the surrounding
atoms to the central atom with
single bond (one shared pair)
4. Determine how many
electron pairs you have left
5. Place lone pairs around
terminal atoms to satisfy octet
(stick any leftovers on central
atom)
6. If the central atom has not
achieved an octet, form
multiple bonds to do so
COCl2
NO3-
Honors Organic Chemistry
Lewis Structures – Additional Practice
H2O
CO32-
NH3
SOCl2
CH2O
C2H4
Resonance Structures
Often, it is possible to draw more that one legal Lewis structure for a molecule.
e.g. SO42-
Expanded Octet:
Each of these structures fulfills the requirements of a Lewis structure, therefore, they are valid
resonance structures. No single one of these resonance structures gives a complete picture of
what the ion looks like. A resonance structure is a lewis structure that contributes to the overall
resonance hybrid of a molecule or polyatomic ion.
Honors Organic Chemistry
Resonance Hybrids – What does a molecule really look like?
All of the resonance structures that can be drawn for a sulfate ion are not created equal – how do
we decide which ones contribute most to the actual picture of what a molecule looks like? To do
this, we have to minimize the formal charge of each atom.
CO2
By minimizing the formal charge on each atom, we can isolate the resonance structure that best
represents what a molecule or ion looks like.
Try some others:
C2O42-
SO32-
General Guidelines for Resonance Structures
1.
2.
3.
4.
5.
6.
7.
8.
Try to draw structures that are as low in energy as possible
The best structures tend to have the maximum number of bonds and the most octets
When structures are equivalent in terms of bonds and octets, minimize formal charge
to find the more stable structure
All structures must be valid. Only electrons may be moved to change between
structures, bonding sequence of atoms must remain the same.
Use curved arrows to show the movement of electrons. Only move lone pairs and
multiple bonds.
Separate resonance structures by a double headed arrow.
Resonance stabilization is very important when it delocalizes or spreads a charge over
two or more atoms
Negative formal charges are more stable on atoms with higher electronegativities.
Honors Organic Chemistry
Worksheet – Resonance and Formal Charge
1.
For each of the following compounds, draw the important resonance structures.
Indicate which structures are major and minor contributors or whether they have the
same energy
a. [H2CNO2]e. [H2CCN]-
:O:
||
b. H – C – NH2
:O:
(-)
.. || ..
f. H2C – C – O – CH3
..
c. [H2COH]+
:O:
:O:
|| ..
||
g. H – C – CH – C – H
d. H2CNN
2.
Draw the important resonance structures for the following molecules and ions
a. H2C = CH – CH2+
b. H2C = CH – CH2-
Honors Organic Chemistry
Summary of Bonding Patterns
Atom
Valence
Electrons
Positively
Charged
Neutral
B
C
N
O
Halogen
Important Note: The bonds shown do not always have to be single bonds
Ex:
Negatively
Charged
Honors Organic Chemistry
VSEPR Theory
VSEPR:
VSEPR theory is used to predict the 3-D geometry of the terminal atoms around the central atom
of a molecule – VSEPR theory tells us the shape of the molecule.
How to determine the 3-D geometry of a molecule from a Lewis Structure
Ex: H2O
1. Count the number of regions of electron
density around the central atom
This tells you the parent shape of the molecule
(see below)
2. Use the number of bonded regions of
density and the total regions of density to
determine the actual shape of the molecule
Ex: SCl4
3. Remember that lone pairs repel more
strongly than bonded pairs, so in a trigonal
bipyramid, place them on the plane as opposed
to the axis
Predict the parent shape and the 3-D geometry of the following molecules:
NH3
SiO2
IF3
IF5
Honors Organic Chemistry
Parent
Shape
Linear
Trigonal
Planar
Tetrahedron
Trigonal
Bipyramid
Octahedron
Pentagonal
Bipyramid
Honors Organic Chemistry
Hybridization
Consider the molecule methane, CH4. One carbon atom bonded to four different hydrogen
atoms.
Electron Configurations:
C:
___
1s
___
2s
___ ___ ___
2p
H:
___
1s
Hund’s Rule:
Pauli Exclusion Principle:
We know that single bonds are formed by the overlap of two orbitals that each have one unpaired
electron – to be able to form 4 bonds to carbon, one electron will have to be promoted from the
2s to the empty 2p orbital.
C:
___
1s
___
2s
___ ___ ___
2p
At present, of the four bonds that carbon is making to hydrogen, there are three “1s overlapping
2p” bonds and one “1s overlapping 2s” bond. From your knowledge of s and p orbitals, draw a
diagram of what this molecule would look like.
There are a couple of problems with the diagram above.
First, experimental evidence shows us that the bond angles for the above molecule are
incorrect – all 4 hydrogen atoms should be equivalently spaced 109.5° from one
another.
Second, all experimental evidence tells us that the four bonds in methane are equal in
energy.
To take into consideration these two problems, a bonding theory called hybridization was
devised
Hybridization:
Honors Organic Chemistry
Hybridization Continued
sp hybrids:
produced from one “s” and one “p” orbital. Two “p” orbitals remain unchanged and
available for π bonding.
Ex: Beryllium in BeCl2 – two sp hybrid orbitals form two σ bonds.
___ ___ ___
2p
___
2s
Electron
Promotion

___ ___ ___
2p
___
2s
Drawing
___ ___
Hybridization
2p

___ ___
sp
Ex: Carbon in C2H2 – two sp hybrid orbitals form two σ bonds. Two “p” orbitals remain available for π
bonding.
Drawing
___ ___ ___
2p
___
2s
Electron
Promotion

___ ___ ___
2p
___
2s
___ ___
Hybridization
2p

___ ___
sp
sp2 hybrids: produced from one “s” and two “p” orbitals.
One “p” orbital remains unchanged and
available for π bonding.
Ex: Boron in BeF3 – three sp2 hybrid orbitals form three σ bonds.
___ ___ ___
2p
___
2s
Electron
Promotion

___ ___ ___
2p
___
2s
Hybridization

Drawing
___
2p
___ ___ ___
sp2
Ex: Carbon in C2H4 – three sp2 hybrid orbitals form two σ bonds. One “p” orbital remains available for π
bonding.
Drawing
___ ___ ___
2p
___
2s
Electron
Promotion

___ ___ ___
2p
___
2s
Hybridization

___
2p
___ ___ ___
sp2
sp3 hybrids: produced from one “s” and three “p” orbitals.
Ex: Carbon in CH4 – four sp3 hybrid orbitals form four σ bonds.
___ ___ ___
2p
___
2s
Electron
Promotion

___ ___ ___
2p
___
2s
Hybridization

___ ___ ___ ___
sp3
Drawing
Honors Organic Chemistry
sp3 hybrids continued:
Ex: Oxygen in H2O – two sp3 hybrid orbitals with unpaired electrons will form two σ bonds. The two sets
of paired electrons exist as lone pairs.
Drawing
___ ___ ___
2p
___
2s
sp3d hybrids:
There is no electron
Electron
Promotion

promotion because
all orbitals are occupied
Hybridization

___ ___ ___ ___
sp3
produced from one “s”, three “p”, and one “d” orbitals.
Ex: Phosphorus in PCl5 – five sp3d hybrid orbitals form five σ bonds.
___ ___ ___ ___ ___
3d
___ ___ ___
3p
Electron
___
Promotion
3s

___ ___ ___ ___ ___
3d
___ ___ ___
3p
Hybridization
___

3s
___ ___ ___ ___
3d
___ ___ ___ ___ ___
sp3d
Ex: Sulfur in SF4 – five sp3d hybrid orbitals form four σ bonds. One lone pair remains.
___ ___ ___ ___ ___
3d
___ ___ ___
3p
Electron
___
Promotion
3s

___ ___ ___ ___ ___
3d
___ ___ ___
3p
Hybridization
___

3s
___ ___ ___ ___
3d
___ ___ ___ ___ ___
sp3d
Ex: Iodine in IF3 – five sp3d hybrid orbitals form three σ bonds. Two lone pairs remain.
___ ___ ___ ___ ___
3d
___ ___ ___
3p
Electron
___
Promotion
3s

___ ___ ___ ___ ___
3d
___ ___ ___
3p
Hybridization
___

3s
___ ___ ___ ___
3d
___ ___ ___ ___ ___
sp3d
Ex: Xenon in XeF2 – five sp3d hybrid orbitals form two σ bonds. Three lone pairs remain.
___ ___ ___ ___ ___
3d
___ ___ ___
3p
Electron
___
Promotion
3s

___ ___ ___ ___ ___
3d
___ ___ ___
3p
Hybridization
___

3s
___ ___ ___ ___
3d
___ ___ ___ ___ ___
sp3d
Honors Organic Chemistry
Worksheet: Hybridization and Lewis Structures
1.
For each of the following, add lone pairs where needed to form octets, predict the
hybridization and bond angles for the circled atoms.
a)
O
||
H–O–C–O
b)
H H
| |
H–C–C–O
| |
H H
c)
d)
H H
| |
H–C–C
| |
H H
e)
H
|
Cl – C = C – C = N – H
| |
H H
f)
CH3
|
CH3 – N – CH3
|
CH3
H–C≡C–H
a. Hybridization: _____________ Bond Angle: _______
b. Hybridization: _____________ Bond Angle: _______
c. Hybridization: _____________ Bond Angle: _______
d. Hybridization: _____________ Bond Angle: _______
e.
i. Carbon - Hybridization: _____________ Bond Angle: _______
ii. Nitrogen - Hybridization: _____________ Bond Angle: _______
f. Hybridization: _____________ Bond Angle: _______
2.
From question #1, which structures are ions?
3.
What type of orbitals are overlapping between the atoms in the following?
a)
b)
H
| ..
H – C = O:
c)
H
H
|
|
H–C=C=C–H
H
| ..
H – C – O:
|
|
H H
4.
5.
6.
Honors Organic Chemistry
Draw Lewis Structures for the following. Remember, hydrogen and halogens do not
make double bonds.
a. CO2
d. N2H4
g. NO2+
b. SeCl4
e. ICl3
h. BF3
c. NH4+
f. SiO3-2
State the hybridization of each central atom from Question #4
a. ______
d. ______
g. ______
b. ______
e. ______
h. ______
c. ______
f. ______
..
Use wedges and dashes for σ bonds to draw an orbital diagram of CH3 – C – O:
Draw any “p” orbitals, label bond angles.
||
|
:O: H
Honors Organic Chemistry
Worksheet – Bonding and Hybridization
1.
For each of the following compounds:
a. Give the hybridization for each atom except hydrogen
b. Give the approximate bond angles for each atom except hydrogen
c. Draw an orbital diagram using lines, wedges and dashed lines for sigma bonds.
Draw the “p” orbital interaction for pi bonds
I.
H3O+
III.
CH3 – C = N – H
|
H
II.
(CH3)4N+
IV.
CH2O
2.
For each of the following:
a. Draw the Lewis structure
b. Indicate what type of orbitals are overlapping to form each bond
c. Give approximate bond angles for each atom except hydrogen
I.
CH3 – C ≡ C – CHO
III.
(CH3)2NH
II.
H2N – CH2 – CN
IV.
CH3 – CH = C(CH3)2
3.
Predict the hybridization and geometry of each carbon atom in the following anion:
O
(-)
|| ..
CH3 – C – CH2
4.
On a separate sheet of paper, draw orbital diagrams of the pi bonding in the following
compounds. Use lines, dashes and wedges to show sigma bonds
a. CH3COCH3
b. CH3 – C ≡ C – CHO
c. Cis  CH3 – CH = CH – CH2CH3 (circle the six coplanar atoms in this molecule)
Honors Organic Chemistry
Worksheet – Bonding Part 2
1.
2-pentyne has the formula CH3CCCH2CH3. Use dashed lines and wedges to draw a
3-D diagram of this molecule. Draw p orbitals as clouds. Circle the four atoms that
are in a straight line.
2.
Which of the following show geometric isomerism? Draw the cis and trans isomers
of the ones that do.
a. CH2=C(CH3)2
b. CH3CH=CHCH3
c. CH3C≡CCH3
d.
e. CH3CH = C – CH2CH3
|
CH2CH3
Honors Organic Chemistry
Worksheet – Bonding Part 2 (cont.)
3.
State the relationships between the following pairs of structures. Your choices are:
identical compound, geometric isomers, structural isomers, totally different
molecules.
a. CH3CH2CH2CH3 and (CH3)3CH
b. CH2 = CH – CH2Cl and CHCl = CHCH3
c. CH3
CH3
\
/
CH = CH
d. CH3
CH3
\
/
CH = CH
and
CH3
\
CH = CH
\
CH3
and
e.
and
f.
and
g.
and
CH2
||
CH3 – C – CH3
Honors Organic Chemistry
Material Covered on the Unit 1 Test
1.
Be able to define and give examples of the all of the key terms on page 2 of this
packet
2.
Know the structure and names of all functional groups mentioned in page 2 of this
packet
3.
Be able to draw valid Lewis structures for polyatomic ions and organic molecules
4.
Be able to assign formal charges to atoms in Lewis structures and also be able to
determine which structures are the major and minor contributors to the resonance
hybrid
5.
Be able to predict bond angles and molecular geometry (shape) from Lewis structures
6.
Be able to determine the hybridization that carbon, nitrogen, oxygen and sulfur atoms
have undergone in a molecule or ion
7.
Memorize the relative order of electronegativities of the following elements:
F, O, Cl, N, Br, S, I, C, P, H.
8.
Be able to assign polarity to bonds in Lewis structures
9.
Be able to draw Lewis structures and line-angle diagrams for structural isomers if
given a molecular formula
10.
Be proficient in the use of curved arrows to show electron movement when drawing
contributing resonance structures
11.
Be able to write and interpret condensed structural formulas
12.
Draw orbital diagrams of sigma and pi bonding
13.
Be able to draw representations of 3-D molecules by using wedges and dashed lines