Ch. 11 12 Review

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Polarity, IMFs, Solids
Chapter 8 & 9
Chapter 11- section 1 &2
Chapter 12 – section 1
Lewis Structures
• A Lewis structure is a representation of a
molecule that shows BOTH the bonding
between atoms AND how the valence electrons
are arranged among the atoms in the molecule.
• Octet Rule - all atoms (except H, Be, B) try to
get 8 valence electrons around themselves.
– H wants 2 electrons
– Be wants 4 electrons
– B wants 6 electrons
RULES FOR MAKING LEWIS
STRUCTURES
1. Obtain the sum total of the VALENCE electrons
of ALL the atoms in the molecule.
2. For ion structures, add or subtract electrons to
the total as indicated by the charge on the ion.
Examples: -2 ion, add 2 electrons; +1 ion,
subtract 1 electron
3. Determine the central atom
–
–
A single atom is often the central atom
The atom that can form the most bonds is often the central
atom
4. Start by making single bonds
– In most cases, there is a symmetrical
arrangement of the atoms.
5. Arrange the remaining electrons to satisfy the
octet rule. These electrons are called "lone
pairs".
6. If you do not have enough electrons to satisfy
all atoms, then double or possibly triple bonds
will be required.
7. If you have too many electrons the extras make
an “expanded octet” on the central atom.
Valence Shell Electron Pair
Repulsion Theory
(VSEPR)
• Used to determine the SHAPE of a
molecule - shape helps determine the
properties
• ALL electron pairs (bonding and lone
pairs) try to get as far apart as possible all electrons repel each other
Electronegativity
• The measure of an atom’s attraction for a pair of
BONDING electrons
- low electronegativity is characteristic of metals
- high electronegativity is characteristic of nonmetals
- Electronegativity generally decreases going
down a group and increases across a series.
- Fluorine has the highest electronegativity
Electronegativity The bond is:
Difference
Example
(difference)
0-0.4
Nonpolar
Cl-Cl
(0.0)
0.4-1.0
Polar
H-Cl
(0.9)
1.0-2.0
Very polar
H-F
(1.9)
>2.0
Ionic
NaCl
(2.1)
Note:
• You will not necessarily have electronegativity values on the
test.
– Know the general trends on the periodic table
– Know that:
• If there is a metal, it’s ionic
• Any bond with fluorine is polar
• Common polar bonds: H-N, H-O, H-F, H-Cl, C-N, C-Cl,
S-O and C-O
• Common nonpolar bonds: C-H, anything bonded to itself
Molecular Polarity (p.343)
Nonpolar – Molecule contains only nonpolar
bonds OR contains polar bonds that are
arranged symmetrically.
Polar – Molecule must contain at least 1
polar bond and the bonds are not
arranged symmetrically. This causes a
charge distribution within the molecule.
*You must know the shape of the molecule!
Predicting Molecular Polarity
When there are no polar bonds in a molecule,
there is no permanent charge difference
between one part of the molecule and another,
and the molecule is nonpolar.
For example, the Cl2 molecule has no polar
bonds because the electronegativity is identical
for both atoms.
It is therefore a nonpolar molecule.
Cl - Cl
Predicting Molecular Polarity
A molecule can possess polar bonds and still be
nonpolar. If the polar bonds are evenly (or
symmetrically) distributed, the bond dipoles cancel
and do not create a molecular dipole.
For example, the three bonds in a molecule of BF3
are significantly polar, but they are symmetrically
arranged around the central boron atom. No side of
the molecule has more negative or positive charge
than another side, and so the molecule is nonpolar:
Predicting Molecular Polarity
A water molecule is polar because (1) its O-H bonds
are significantly polar, and (2) its bent geometry
makes the distribution of those polar bonds
asymmetrical.
The side of the water molecule containing the more
electronegative oxygen atom is partially negative, and
the side of the molecule containing the less
electronegative hydrogen atoms is partially positive.
+ side of molecule
- side of molecule
EXAMPLES – Predicting Molecular Polarity:
Decide whether the molecules represented by the following
formulas are polar or nonpolar.
a. CO2 b. OF2 c. CCl4 d. CH2Cl2 e. HCN
a.
The C-O bonds are polar.
The Lewis structure and shape for CO2 is
If we put arrows into the sketch for CO2,
this side negative
this side negative
This is a symmetrical arrangement so this molecule is
NONPOLAR
EXAMPLES – Predicting Molecular Polarity:
Decide whether the molecules represented by the following
formulas are polar or nonpolar.
a. CO2 b. OF2 c. CCl4 d. CH2Cl2 e. HCN
b.
The O-F bonds are polar .
The Lewis structure for OF2 is
The molecular geometry of OF2 is bent
and adding arrows gives:
this side this side +
This asymmetrical distribution of polar bonds would
produce a POLAR molecule.
EXAMPLES – Predicting Molecular Polarity:
Decide whether the molecules represented by the following
formulas are polar or nonpolar.
a. CO2 b. OF2 c. CCl4 d. CH2Cl2 e. HCN
c.
The C-Cl bonds are polar.
The Lewis structure for CCl4 is
The molecular geometry is tetrahedral
and adding arrows gives:
this side -
this side -
This is a symmetrical arrangement and produces a
NONPOLAR molecule
EXAMPLES – Predicting Molecular Polarity:
Decide whether the molecules represented by the following
formulas are polar or nonpolar.
a. CO2 b. OF2 c. CCl4 d. CH2Cl2 e. HCN
d.
The C-Cl bonds are polar and the C-H bonds are nonpolar.
The Lewis structure for CH2Cl2 is
The molecular geometry is
tetrahedral and adding
this side +
arrows gives:
this side -
This is an asymmetrical arrangement that produces a
POLAR molecule.
EXAMPLES – Predicting Molecular Polarity:
Decide whether the molecules represented by the following
formulas are polar or nonpolar.
a. CO2 b. OF2 c. CCl4 d. CH2Cl2 e. HCN
e.
The H – C bonds are nonpolar and the C – N bond is polar.
The Lewis structure is:
The molecular geometry is linear
and adding arrows gives:
this side -
this side +
This is an asymmetrical arrangement and the molecule is
POLAR.
Intermolecular Forces (IMF)
• These are forces that exist BETWEEN
particles of a substance (for example: the
forces between 2 hydrogen molecules NOT the bond holding the 2 hydrogen
atoms together to make the molecule)
Intermolecular Forces
Weakest IMFs
Strongest IMFs
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Dispersion Forces
An instantaneous dipole can form when electron density
becomes unbalanced because electrons move as seen
in the helium atom below
Frame 1
Frame 2
Frame 3
When both electrons end up
on one side of the atom (like
in frame 3), that side of the
atom acquires a δ- charge and
the other side becomes δ+
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Dispersion Forces
20
• These forces are present between particles of
ALL substances
IN GENERAL:
• Their relative strength is related to the total number of
protons & electrons in a single particle (a larger molar
mass indicates more protons & electrons - stronger
dispersion forces)
• London dispersion forces increase as surface area
increases. Increasing surface area makes the
molecule more polarizable (able to form
instantaneous dipole easier)
Higher molar mass means
larger dispersion forces and
higher mp and bp
Higher molar mass because
more protons, neutrons and
electrons – more charged
particles – more attraction.
We can use the molar mass
to determine the relative
strength of dispersion forces
BUT the REASON is the
increase in the number of
charged particles NOT the
increase in mass
22
same formula with same molar mass BUT branching
decreases the dispersion forces (decreases the
polarizability of the molecule’s electron cloud) so
branched structure has lower boiling point (bp) 23
Dipole Forces
Only found between molecules that are polar
• Polar molecules form dipoles (regions of partial
charges formed by the unequal sharing of the
electrons in the polar bond)
Dipole forces are stronger
than dispersion forces
Relative strength is related to
the polarity of the molecule
Hydrogen Bonding
• Hydrogen bonding is a
special form of dipole
forces.
• It is the strongest dipole
force
• It only occurs when
molecules have a
hydrogen attached to a
VERY electronegative
element (nitrogen,
oxygen or fluorine)
Examples of H-bonding
Notice that H “bonds” to the lone pair on the O or N
26
To maximize H-bonding,
water molecules actually
get farther apart when
changing from the liquid
state to the solid state.
This makes ice
less dense than water.
27
High BP due
to presence of
H bonding
Showing the strength of hydrogen
bonding using BP data
28
Ion-Dipole Force
• An ion-dipole force is an interaction between an ion (e.g.,
Na+) and the partial charge on the end of a polar molecule
(e.g., water).
• Not really an intermolecular force because there is more
than 1 substance involved
• important for aqueous solutions of ionic compounds
• the strength of the ion-dipole attraction is one of the main
factors that determines the solubility of ionic compounds in
water
• Example: NaCl (aq)
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Ionic Bonds
• Are only found between positive and
negative ions in an ionic compound
IMF Flow Chart
Comparing IMFs
• Dispersion forces are found in all substances.
– Their strength depends on molecular shapes and the number of
charged particles (molar mass).
• Dipole-dipole forces add to the effect of dispersion forces.
– They are found only in polar molecules.
– H-bonding is a stronger special case of dipole-dipole interactions.
H-bonding is only significant when H is bonded directly to
N, O or F.
• Ion-dipole and ionic bonding are possible when there are ions.
– Ion-dipole interactions are slightly stronger than H-bonds.
– Ionic bonding attractions are stronger than ion-dipole forces.
IN GENERAL
Ionic bonding > Ion-dipole > H-bonding > dipole-dipole > dispersion
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Solids
Types of solids:
• Ionic
• All ionic compounds are solids
• Ions held by the attraction between cations and
anions (ionic bond)
• Generally brittle, not electrical conductors, held in
crystal structures
• Metallic Solids
– Held together by a “sea” of electrons
– Electrical conductors
– Alloys (p.473)
• Mixtures of metals
• Bronze = copper + tin
• Metals are mixed to form materials with desirable
properties.
• Covalent-Network Solids
– Held together by covalent bonds
– Often very hard (diamond)
• Polymers
– long chains of atoms (often carbon)
• Molecular Solids
– Held together by strong IMF
– Usually softer than other types of solids

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