Laboratory II: Modeling
Molecular Structure Part I
Purpose: Learn about the various
structure and chemical bonding patterns
of a variety of compounds. Specifically:
Lewis structures, hybrid orbital's,
VSEPR, bond/molecular polarity,
resonance and chemical nomenclature.
Laboratory II
Commonality: Each model HELPS to represent and
understand characteristics of molecules.
• Model 1-Ball and Stick: Sense of the way atoms are
arrayed in molecules and of the molecule's threedimensional structure
– Why is 3D structure important?
• Molecular geometry: the three-dimensional arrangement of the
atoms that constitute a molecule.
– Molecular Geo: helps determine several properties of a substance
including its reactivity, polarity, phase of matter, color, magnetism,
and biological activity.
Laboratory II
• Issue? Why not stop with Ball and Stick?
– It is essential for a model to extend beyond a
simple appreciation of a static object but give us
some sense of the object's potential.
• This model building exercise is designed to show you
the value of molecular modeling structures have and
to better inform you on molecular behavior and the
limits of each model.
Model II: Lewis Dot StructureRepresentation using Lewis symbols of covalent
bonding in a molecule
• Diagrams that show the bonding
between atoms of a molecule, and
the lone pairs of electrons
associated with the molecule.
– Shared electron pairs are shown as
lines
– Unshared electrons are show as lone
or paired dots
• ONLY valence-shell electrons are
shown
• Lewis Model Uses Rule of
Orbital's, Octet Rule, and the
Principle of Electroneutrality
The Rule of Orbital's: the total number
of lone pairs and bond pairs (LP+BP)
associated with an atom cannot
exceed the number of Valence Shell
Orbital's (VSO)
• VSO = n2, where n is the row of the
Periodic Table in which that atom resides
Examples:
Hydrogen- n = 1: max VSE pairs (LP+BP) =
VSO = 1
(B, C, N, O, F)- n = 2 : max VSE pairs
(LP+BP) = VSO = 4 ("octet rule")
Model II- Lewis Dot Structure
Octet Rule: Atoms tend to bond in Principle of Electroneutrality: Each
such a way that the atoms posses
or share a total of 8 valence shell
electrons
– Gives atoms the same electronic
configuration as a noble gas.
– s2 p 6
• The quantum theory is used to
explain the valence electron
energy levels.
– Molecules tend to be most stable when
the outermost electron shells of their
constituent atoms contain eight
electrons.
– Explains the stability and un-reactivity
of noble gases (except helium)
atom in a covalent molecular
assembly has a formal charge as close
to zero as possible.
• Formal Charge (FC): The Charge of any
atom in a molecule is the charge the atom
would have if all the constituent atoms had the
same electronegativity.
Rules and Theories help explain
placement and rearrangement of
electrons and formation of bonds.
Laboratory II: Steps for Lewis Structure
1.) Sum the Valence electrons.
•
•
If you have any cations (+) subtract them
from the total VE (group #)
If you have anions (-) add them to the total
VE.
2.) Write the central atom then the
attached atoms.
•
•
The central atom is generally the least
electronegative atom (not including
hydrogen).
Electronegativity increases in the periodic
table as you go from left to right and as
you go up.
EXAMPLE: HCN
Step 1: H (1), C(4), N (5)
1+4+5-0 = 10 VSE  5 VSE pairs
Step 2a:
Step 2b:
C
Laboratory II: Steps for Lewis Structure
Cont.
3.) Complete the octets around the
atoms bonded to the central atom.
•
EXAMPLE: HCN
Step 3: The two bonds use up four
electrons so 6e- remaining
Make ‘Happy atoms’
a.) Start w/minimal bonds
b.) Complete Octets on atoms around
C-atom
Laboratory II: Steps for Lewis Structure
4.) Place any and all leftover atoms
on the central atom, even if doing
so violates the ‘normal’ octet rule.
• If there are not enough electrons
to give the central atom an octet,
borrow e- from the adjacent
molecules to form multiple
bonds
Example HCN
• The C-atom has 4e-, but according
to Step 4, needs 8e– What to do? Rearrange e- to
make bonds.
Rearrange e-
Carbon Atom is still not Happy! Make
more bonds!
Happy Molecules!
Laboratory II: Important Concepts of
Lewis Structures
Most of the atoms in this
exercise obey the “octet
rule”
Exceptions and Violation of Octet
Rule- When to Expand•
• Atoms with an expanded octet can
•
accommodate more than eight
electrons.
– Atoms in the third and higher
periods of the periodic table
– Transition metals
– Representative groups (IIIAVIIA)
o Initially, assume that all heavy
atoms obey the octet rule
o We will discover exceptions.
•
•
Molecules and polyatomic ions
containing an odd number of electrons
Molecules and polyatomic ions in which
an atom has fewer than an octet of
valence electrons
Molecules and polyatomic ions in which
an atom has more than an octet of
valence electrons
The number of valence electrons is odd
so a complete pairing of these electrons
is impossible, in other words an octet
around each atom is not possible.
Laboratory II Procedure
• Build ‘Ball and Stick’ of 16 molecules and deduce name,
Lewis structures, geometry, polarity, resonance, and possible
isomerism of each molecule.
• Work in pairs to build molecules together; DO NOT copy!
– Work together to determine the answers
• The models you will use consist of wooden balls (atoms), pegs
(single bonds) and springs (used for multiple bonds).
– The balls represent the nuclei and inner core electrons.
– The pegs and springs represent valence shell bonding or nonbonding
(lone pair) electron pairs.
Laboratory II: Concepts to Note
Difference between Resonance and an Isomer
Resonance: Individual Lewis
Structures in cases where 2 or
more Lewis Structures are
relatively equally adequate
descriptions of the same
Molecule.
• Structures must be stable- can not
break rules
• The resonance structures in such
instance are averaged to give a
correct description of the real
molecule.
– Blend of all structures
What does a Unicorn, Dragon, and Gryphon all have
in common?
Resonance
They are NOT REAL depictions
of the molecule.
• The real structure is a blend,
compilation, an average, some
intermediate of all the resonance
structures
– Do not get caught up on static
representation
• Just moving electrons to make or
break bonds!
Resonance structures help explain
molecule’s stability and/or
reactivity
Example: Benzene Ring
Resonance
Be Reasonable
• Follow the octet and
expanded octet rules
• Make molecules with
Formal Charges close to
zero as possible
• Satisfy VSERP
BAD Representation!!!!!
** Isolated Formal Charges
** Octet not satisfied for Nitrogen atom
Structural Isomer
Isomer: Compounds whose
molecules have the same overall
composition but DIFFERENT
STRUCTURES
•
If it is possible to arrange the atoms
of your molecule in another skeleton
structure that is different from the one
already made than it is an isomer
– Same molecular formula but
different structural formula
Molecular formula: Chemical
formula-describes what atoms are
present in the molecule
Example: Formaldehyde- CH2O
Structural formula: graphical
representation of the molecular
structure, showing how the atoms are
arranged
Formaldehyde
Work an Example Together
Lab Example: Formaldehyde CH2O
Step 1: Draw Lewis Dot Structure
Step 2: Examine the Molecule's Molecular
–
–
Geometry and describe the general
overall shape of the molecule.
Atoms linear? Triangle Shape?
Use your textbook and black binder to
help with geometry
Formaldehyde is flat (all of its
atoms lie in one plane) with
its peripheral atoms forming
the vertices of a triangle.
Structure is called trigonal planar
Step 3: Predict the Molecule's Polarity
Polar bonds are formed when two
atoms with different
electronegativity bond.
The more electronegative atom will get more than its
fair share of electrons, and be the negative end of
the bond.
The less electronegative atom will have some electron
density pulled away, so it will be the positive end of
the dipole.
Step 3: Predict the Molecule's Polarity
Total dipole moment of the
molecule is the vector sum
of the bond dipoles.
– If dipoles oppose one another
they will cancel out
From the shape and EN predict
whether the molecule as a
whole should be polar.
– If you believe it to be polar,
indicate the direction of the
molecule's dipole moment on
the Lewis Dot Formula you
drew for formaldehyde
Formaldehyde: The difference
in electronegativity between C, at
2.5, and H, at 2.1, is not
sufficient to give a significant
bond dipole.
The difference between C, at 2.5, and
O, at 3.5, is markedly polar.
Step 4. Examine the Molecule's
Potential for Resonance
• There is only one place
for the double bond in
formaldehyde.
• There is no place for a
triple bond to form
– Conclusion: No stable
resonance structures.
Step 5. Examine the Molecule's
Potential for Structural Isomerism
• If your molecule has structural
isomers, build a model of at least
one of them and compare it with
your original model.
• It is possible to arrange the atoms
of formaldehyde
– Build Molecule
– Deduce Formal Charge
– Stable?
Methanal: chemical formula: CH2 O
Formal Charges:
Valence e- - (1/2 bonding e- + nonbonding e-)
Oxygen: 6- (1/2 *8 +0) = +2
Carbon: 4- (1/2 *4 +4)= -2
Due to separation of formal charges
on the O and C atoms, there is no
isomer
Molecules to Explore
a. Methane (CH4):
–
–
–
–
Construct a ball and stick model
• Note the symmetry
Discern saturated or unsaturated: Saturated
Note Geometry: tetrahedral molecule
Carry out the additional structural analyses
described in Part B above, like we did for
formaldehyde
b. Dichloromethane (CH2Cl2) :
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–
–
–
Convert to CH2Cl2 by replacing two
hydrogen atoms with chlorine atoms.
Do the additional structural analyses
described in Part B above.
Draw the Lewis dot structures for all of the
structural isomers of CH2Cl2 that you think
may exist.
Build structures of all of your isomers
• Compare them; are they superimposable?
Will not adequately describe symmetry
Better representation
Molecules to Explore Cont.
c. Ethyl alcohol (ethanol): CH3CH2OH)
d. Dimethyl ether: (CH3OCH3)
•
Ethanol Key Points:
Very soluble in water
Bp: 78.5°C
Construct models for ethyl alcohol and
dimethyl ether.
•
•
These two compounds are structural
isomers, but they have very different
properties.
The differences are caused by the presence
of the O-H bond.
 The hydrogen bonding orientation is
an important contribution to the forces
between molecules in ethanol.
Dimethyl ether Key Points:
Not water soluble
Bp: -24°C
Assignments Due Date
• What is due: the worksheet
with 16 molecules
completely answered
• When is it Due: Hand into
me by 4pm tomorrow
• Grade: based on correct
answers, completeness, and
turned into me on-time.
• Notebook Check: make
sure to put any example
molecules in your
notebook
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Formaldehyde
Methane
Dichloromethane
Ethyl alcohol