Chapter 10 Shapes of Molecules

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Chapter 10Shapes of Molecules
The Three Dimensional Reality of Molecular Shapes
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The Three Dimensional Reality of Molecular Shapes
Each atom, bonding e- pair, and lone e- pair has
its own position in 3D space. Position
determined by the attractive and repulsive forces
that govern all matter.
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Molecular Shape is Crucial to Life Processes
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Molecular Shape is Important in
Nanotechnology Nanotechnology Building tiny
machines at the molecular level.
Buckyball (fullerenes) a C60 structure
discovered in soot in 1985 Many, many uses One
use may be to sneak medicines into cells or
through the blood brain barrier, where some
substances (medicines) may not normally enter.
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Section 10.1 Depicting Molecules and Ions with
Lewis Dot Structures 2D before 3D
Lewis dot structures consist of two parts
(1) Element symbol nucleus inner
electrons Ex The element lithium has an element
symbol Li
(2) Surrounding dots valence electrons (outer
most shell)
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Section 10.1 Depicting Molecules and Ions with
Lewis Dot Structures
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Continued

• Steps for writing Lewis Structures for molecules
  that have only single bonds
• Determine the total number of valence e-
  available
• Example NF3 N ? 5 e- F ? 7 e-
  (x 3) 21 e- Total 26 e-

• Distribute the remaining e- in pairs so that each
  atom has 8 e- (except H, has 2 e-)
• 1st Place lone e- pairs on surrounding (more
  EN) atoms to give each full valence.
• 2nd If any e- remain, place them on the central
  atom.
• Example NF3

(5) Check that each has a full valence shell.
DONE.
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Section 10.1 Depicting Molecules and Ions with
Lewis Dot Structures
Lewis Dot Structures are not 3D, so several
different depictions are correct.
Etc.
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Section 10.1 Depicting Molecules and Ions with
Lewis Dot Structures
In cases of multiple bonds (double, triple),
there is an additional step (6) If a central
atom still does not have an octet (after Step 4),
make a multiple bond by changing a lone pair
from one of the surrounding atoms into a bonding
pair to the central atom. Examples
CH4 N2
In cases where there is a polyatomic ion, Lewis
Dot Structure is shown in square brackets
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Section 10.1 Resonance Structures Have the
same relative placement of atoms, but different
locations of bonding and lone e- pairs.
Example Nitrite (NO2-)
Double bonds change location. Lone pair on O atom
changes location.
Neither structure depicts nitrite accurately
The two O,N bonds in this compound have bond
lengths and bond energies that lie somewhere
between the ON and the ON bond.
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Section 10.1 Resonance Hybrids Resonance
structures are not real bonding depictions
Structure I Structure II
In reality, Structure I and Structure II
do not switch back and forth from one instant
to the next. The actual nitrite molecule is an
average of the two. The e-s are delocalized
over the entire molecule. (Just as e-s in a
metallic bond are delocalized around the entire
sea of electrons.)
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Section 10.1 Resonance Hybrids Sometimes
implied (no indication). Sometimes indicated
by dotted lines.
Example Ozone (O3)
Example Benzene (C6H6)
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Draw a Lewis Dot Structure for CCl4 CCl2F2 CH4O
NH3O C2H6O (no OH bonds) 10.6 10.8
Resonance Structures 10.10 10.12
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Section 10.1 Which is the more important
resonance structure?
All resonance forms contribute equally when
central atom has surrounding atoms that are all
the same.
When atoms surrounding the central atom are not
the same One resonance form may weight the
average (In other words, it counts more than
the other forms.) Determining the most important
resonance form Determine each atoms formal
charge the charge it would have if the bonding
electrons were equally shared,
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Section 10.1 Which is the more important
resonance structure? Formal charge Determined
for each atom in a compound
valence e- - ( unshared valance e- ½
shared valence e-)
Example Ozone (O3) Formal charge for OA 6
(4 ½4) 0 Formal charge for OB 6 (2
½6) 1 Formal charge for OC 6 (6 ½2)
-1 In the case of ozone, both resonance
structures (I and II) have the same
formal charges (but on different O atoms) so they
contribute equally to the resonance hybrid.
Formal charges must sum to the total charge on
the chemical species.
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Section 10.1 Which is the more important
resonance structure? Formal charge Determined
for each atom in a compound
Recall In nearly all compounds H atoms form
1 bond C forms 4 bonds N forms 3 bonds O
forms 2 bonds Halogens (F, Cl, Br, I,) form 1
bond
When formal charge is 0, an atom has its usual
of bonds.
Ozone Example Formal charge for OA 6 (4
½4) 0 Formal charge for OB 6 (2 ½6)
1 Formal charge for OC 6 (6 ½2) -1
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Section 10.1 Which is the more important
resonance structure? What about a case where the
resonance structures do not contribute equally?
Occurs when there are different atoms around the
central atom.
Example Cyanate ion, NCO-
Criteria for choosing the more important
resonance structure (1) Smaller formal
charges ( or -) are preferable to larger
ones. Resonance form I is out. (2) The
same nonzero formal charges on adjacent atoms are
not preferred. Not applicable to this
example. (3) A more negative formal charge
should reside on a more electronegative atom. O
is more EN than N, so Resonance form III is the
winner.
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Section 10.1 Formal charge is not the same as
the ON
What is the difference?
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Section 10.1 Exceptions (Limitations) to the
Octet and Formal Charge Rules
(1) e- deficient molecules gaseous molecules
containing Be or B as central atom
Halogens much more EN than Be or B Formal
charge rules make sharing of extra lone pairs by
halogens unlikely
Be 4 e- B 6 e-
(2) Odd e- molecules central atom has odd of
valence e- free radicals very reactive (b/c
very unstable) often react with each other to
pair up their lone e- (make you age)
(3) Expanded valence shells more than 8 valence
e- occurs only where d orbitals are available ?
Row 3 or higher (Review of orbital
types p289 295 s 2 e-, p 6 e-, d 10 e-,
f 14 e-)
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Resonance Structures 10.14a 10.16b Other
suggested problems 10.15 10.17 10.19 10.24
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Formal charge valence e- - ( unshared
valance e- ½ shared valence e-)
Criteria for choosing the more important
resonance structure (1) Smaller formal
charges ( or -) are preferable to larger
ones. Resonance form I is out. (2) The
same nonzero formal charges on adjacent atoms are
not preferred. Not applicable to this
example. (3) A more negative formal charge
should reside on a more electronegative atom. O
is more EN than N, so Resonance form III is the
winner.
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Section 10.2 Valence-shell electron-pair
repulsion (VSEPR) theory Molecular shape is
important in many, many scientific disciplines.
Medicine Receptors
Nanotechnology Membrane Transport
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Section 10.2 Valence-shell electron-pair
repulsion (VSEPR) theory Molecular shape is
important in many, many scientific disciplines.
Ecology Talking trees
Jack Schultz, Chemical Ecologist
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Section 10.2 Valence-shell electron-pair
repulsion (VSEPR) theory
Lewis Dot Structures, 2D (Blueprint)
VSEPR, 3D (House)
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Section 10.2 Valence-shell electron-pair
repulsion (VSEPR) theory Each group of valence
electrons around a central atom is located as far
away as possible from the others in order to
minimize repulsions.
e- group can be a single bond, double bond,
triple bond, lone pair, lone e-
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Section 10.3 Molecular Shape and Molecular
Polarity In molecules with more than 2 atoms
Shape and bond polarity determine molecular
polarity.
Molecules with only 2 atoms.
Relative electronegativities of the two atoms
determine polarity.
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Section 10.3 Molecular Shape and Molecular
Polarity In molecules with more than 2 atoms
Dipole moments
Dipole moments a measure of molecular
polarity magnitude of partial charges on ends of
a molecule (in coulombs) x distance
between them
Behavior of Molecules With and Without Dipole
Moments Electric field Polar molecules (which
have a dipole moment) orient with partial
charges towards oppositely charged electric
plates. Molecules with out a dipole moment will
not orient themselves in any particular
direction. Molecules with no dipole moment can
be polar.
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Section 10.3 Molecular Shape and Molecular
Polarity Dipole moments When molecular shape
influences polarity
Large ?EN between C (EN 2.5) and O (EN 3.5) ?
C O bonds are polar CO2 molecule is linear ?
Two identical bond polarities are counterbalanced
(in other words, they cancel each other out) As
a result, CO2 has no net dipole moment.
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Section 10.3 Molecular Shape and Molecular
Polarity Dipole moments When molecular shape
influences polarity
H2O (like CO2) also has two identical molecules
bonded to the central atom.
However, H2O (unlike CO2) has a dipole
moment. Bond polarities are not canceled out
because of the shape of the water
molecule V-shaped rather than linear. The O
end of the molecule is more negative than the H
ends
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Section 10.3 Molecular Shape and Molecular
Polarity When different molecules have the same
shape, the nature of the atoms Surrounding the
central atom can have a major effect on polarity.
CCl4 does not have a dipole CHCl3 has a
dipole
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Effect of Molecular Polarity on Behavior
Example Boiling point of NH3 versus PH3
Why is NH3 boiling point higher?
Also determines reaction behavior NH3 H ?
NH4 (p392)
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A closer look at molecular shapes double bonds
and lone pairs
Bond Angles Idealized vs. Actual
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A closer look at molecular shapes double bonds
Effect of double bonds on bond angle when
surrounding atoms are different.
Rule The double bond, with its greater e-
density, repels the two single bonds more
strongly than they repel each other.
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A closer look at molecular shapes lone pairs
Effect of lone pairs on bond angle.
Rule Lone pairs repel bonding pairs more
strongly than bonding pairs repel each other.
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A closer look at molecular shapes a few more
details
Bond angles for Equatorial groups 90º Axial
groups 120º
General Rule The greater the bond angle, the
weaker the repulsion.
In this case Equatorial-equatorial repulsions
are weaker than axial- rquatorial repulsions.
Implication Lone pairs, which exert stronger
repulsions, will tend to occupy equatorial
positions.
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A closer look at molecular shapes a few more
details
General Rule The greater the bond angle, the
weaker the repulsion.
Implication If two lone pairs are present, they
will always occupy opposite vertices (furthest
apart)
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More VSEPR Practice
(1) Lewis dot structure (dominant resonance form
calculate formal charge if needed) (2) 3-D
geometry (Table 10.9) (3) Molecular polarity
Bond dipoles cancel? Lone pairs present?
Different surrounding atoms?
GeH2 PCl5 SF4 ClF3 XeF2 SF6 BrF5 XeF4 NH4 SO4-
Additional Optional Homework Problems 10.30,
10.35, 10.38, 10.39, 10.51, 10.53, 10.54, 10.97
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Chemical Reaction vs. Physical Interactions
(Chapter 12)
Boiling point of NH3 versus PH3 Why is NH3
boiling point higher? The N H bonds between
NH3 molecules matter, not the N H bonds within
the NH3 molecule.
EN is inversely proportion to atomic radius.
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Summary So Far Overall Polarity of a Molecule
If a molecule is polar, it will (1) have a net
dipole moment (2) orient itself in an electric
field If a molecule is nonpolar, it will (1)
not have a net dipole moment (2) will move about
randomly in an electric field

• Steps used to determine molecular polarity
• Draw the 2-D Lewis dot structure to determine the
• number and types (single, double, triple) of
  bonds
• present, and any lone pairs present.
• When dealing with several options (resonance
  structures),
• determine the dominant resonant structure by
• 1. calculating formal charge for the atoms of
  each
• molecular possibility that you are
  evaluating
• 2. using the three criteria for selecting the
  dominant
• resonance structure based on formal charge
• (2) Determine the 3-D shape of the molecule
• (3) Determine the overall polarity of the molecule

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In addition to the wavelength of energy
interaction with the molecule The symmetry of the
molecule (Lewis dot structure VSEPR!!!) will
also determine whether a photon can be absorbed.
Symmetry is NOT a net dipole moment
Symmetrical mirror image Asymmetrical not
mirror image