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Ionic Bonds
An ionic bond is simply the electrostatic attraction
between opposite charges.
Ions with charges
Q1 and Q2:
Q2
Q1
d
The potential energy is given by:
E
QQ
1
2
d
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Molecular Compounds
The simplest molecule is H2:
Increased electron density draws
nuclei together
The pair of shared electrons constitutes a
covalent bond.
• For individual atoms:
– Chemical symbol represents nucleus and core e-.
– Dots around the symbol represent valence e-.
– Get the # of valence e- from the periodic chart
Al •
•
•
•
••
• Se
•
•
••
• Bi •
•
••
• Sb •
•
••
I
••
•
••
Ar
••
••
•
••
P•
••
•
••
••
•N•
•
•
•Si •
•
••
• As •
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The Octet Rule
Atoms tend to gain, lose, or share electrons
until they have eight valence electrons.
Hydrogen is an exception.
H shares only one
electron to fill its 1s shell
H : H
Most other Group A
elements want an octet:
Ionic and Molecular Compounds
•Formation of sodium chloride:
Na
+
Cl
Na+
[ Cl ]
• Formation of hydrogen chloride:
H
+
Cl
H Cl
A metal and a nonmetal transfer electrons to
form an ionic compound. Two nonmetals share
electrons to form a molecular compound.
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Lewis Structures for Ionic
Compounds
••
• Cl
••
2+ ••
••
2-
••
••
O
2+
Mg
••
• Cl
••
••
2 Cl
••
-
••
•
Mg •
Ba
••
MgCl2
••
• O•
••
••
Ba •
••
•
BaO
Lewis Structures
• Lewis structures are representations of molecules
showing all electrons, bonding and nonbonding.
• A valid Lewis structure for most covalent molecules
should have an octet for each atom except
hydrogen.
H2 is H + H
H H
Cl2 is Cl + Cl
Nonbonding electrons
or H H
Bonding
electrons
Cl Cl
or
Cl
Cl
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Drawing Lewis Structures
1. Add up the valence electrons from all atoms. Add 1
for each - charge and subtract 1 for each + charge.
2. Draw a skeleton structure with single bonds.
3. Complete the octets of atoms bound to central atom.
4. Place extra electrons on the central atom.
5. If the central atom doesn’t have an octet, try forming
multiple bonds.
6. Check formal charges to ensure the best structure
More on this later
:
:
:
:
Drawing the skeleton: H2C(OH)2
• Add up the valence electrons : (4 x 1) + (1 x 4) + (2 x 6) =20
• Identify central and terminal atoms
– Carbon atoms are always central atoms
– Hydrogen atoms are always terminal atoms.
• Central atom → atom w/lowest electronegativity.
• Draw skeleton structure, then distribute valence electrons
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Try one: NF3
Molecular
formula
Sum of
valence e-
:
:F:
:F:
:
Draw
skeleton
:
NF3
N
N s2p3
F s2p5 7e-
Total
X 3 = 21e-
26e-
:
:F:
Remaining
valence e-
5e-
Zero: NF3 is uncharged
Lewis
structure
And another!
Write a Lewis structure for CCl2F2
SOLUTION:
Cl
Cl
C
F
F
:
Step 1: Add up the valence electrons:
4 + (2 x 7) + (2 x 7) = 32
Step 2: Identify central atom and place other
atoms around it.
C
F :
: F:
:
: Cl
:
:
:
: Cl :
:
Steps 3-4: Draw bonds and fill in remaining
valence electrons placing 8e- around each
atom.
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Lewis Structures
Draw Lewis structures for:
HF 8e-s
H F
H2O 8e-s
H O H
or
H
NH3 8e-s H N H
H
or
H N H
H
CH4 8e-s
or
H
H C H
H
or
H
F
O H
H
H C H
H
Double and Triple Bonds
• Atoms can share four electrons to form a double bond
or six electrons to form a triple bond.
O2:
O =O
N2:
N N
• The number of electron pairs is the bond order.
• Double bonds are shorter and stronger than singles
• Triple bonds are shorter and stronger than doubles
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Multiple Covalent Bonds
Draw the Lewis structure for N2
••
•
•
•
•
N N
••
••
••
N N
•
••
••
•
•
N N
••
•
•N
•
••
•
N•
Count valence eDraw skeleton
Distribute electrons
Try multiple bonds to give all atoms an octet
Multiple Covalent Bonds
Draw the Lewis structure for CO2
•
•
••
•
•
••
••
O C O
••
•
•
••
•
•
O C O
••
••
••
••
•
•O
••
••
••
••
••
••
•
• C•
•
••
•
O•
O C O
Count valence eDraw skeleton
Distribute electrons
Try multiple bonds to give all atoms an octet
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more than 1 central atom
PROBLEM:
Write the Lewis structure for methanol
(molecular formula CH3OH)
H
:
SOLUTION:
C
O
H
:
H
H
There are 4(1) + 4 + 6 = 14 valence e-.
Hydrogen can have only one bond so C and O must be
next to each other with H filling in the bonds.
Draw skeleton: C has 4 bonds and O has 2.
Distribute electrons: O has 2 pair of nonbonding e-.
Molecules with Multiple Bonds
PROBLEM: Write Lewis structures for the following:
(a) Ethylene (C2H4)
(b) Nitrogen (N2)
PLAN:
If a central atom does not have an octet, then e- can be
moved in to form a multiple bond.
There are 2(4) + 4(1) = 12 valence e-. H can
have only one bond per atom.
SOLUTION: (a)
:
H
H
C
H
C
H
H
H
H
C
C
H
(b) N2 has 2(5) = 10 valence e-. A triple bond is required to make
octet around each N.
N
.
:
N
.
:
:
: .
N
.
N
N
:
.:
N
.
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Optimizing the structure
Formal charge is just a way of keeping track of the
electrons assigned to each atom
Formal = number of valence Charge
electrons
unshared
electrons
+
1/2 of shared
electrons
• Easier to figure it out visually: Count the number of unshared
electrons and the number of bonds around an atom
• Compare to # of valence e-s: The difference is formal charge
Writing Lewis Structures
• Based on formal charges, the best Lewis structure:
– is the one with the fewest charges (or lowest
overall number of charges).
– puts a negative charge on the most electronegative
atom (like O or F).
Cyanate ion
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Lewis Structures: Practice
• NClO
• H3PO4
• H3BO3
• SO3-2
• NO2-1
• P 2H4
Lewis Structures: Practice
• NClO
• H3PO4
••
•O
•
18 e-
• H3BO3
24 e-
H
••
O
••
• NO2-1
18 e-
••
•O
•
••
N
••
Cl ••
••
32 e-
H
••
O
••
• SO3-2
••
•O
•
H
••
O
••
B
••
N
••
•O
•
••
26 e-
H
• P 2H4
••
O ••
••
14 e-
H
••
•O•
• •
P
•O
•
••
••
•O•
• •
S
••
H
H
P
••
P
••
••
O
••
H
H
••
O•
•• •
H
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Resonance
This is the Lewis
structure we would
draw for ozone, O3.
-1/2
-1/2
+
1.48 Å
1.21 Å
-
• But, in the observed
structure of ozone:
– both O-O bonds are the
same length (1.278 Å).
– both outer oxygens
have a charge of -1/2.
Resonance
• One Lewis structure cannot accurately depict a
molecule like ozone.
• We use multiple structures, resonance structures, to
describe the molecule.
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Resonance
•
•
Formate ion (from formic acid)
• The electrons of the double bond do not sit between
the C and the O, but rather can move among the two
oxygens and the carbon.
• They are not localized; they are delocalized.
Ring Resonance
• The organic compound
benzene, C6H6, has two
resonance structures.
• It is commonly depicted
as a hexagon with a
circle inside to signify the
delocalized electrons in
the ring.
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Exceptions to the Octet Rule
• There are three types of ions or molecules that do not
follow the octet rule:
– Ions or molecules with an odd number of electrons
– Ions or molecules with less than an octet
– Ions or molecules with more than eight valence electrons
(an expanded octet)
Odd Number of Electrons
Though relatively rare and usually quite
unstable and reactive, there are ions and
molecules with an odd number of electrons.
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Fewer Than Eight Electrons
NO!
NO!
NO!
• Consider BF3; giving boron a filled octet would put:
→ negative charge on B
→ positive charge on F
– None of these would be very accurate
• Instead, the Lewis structure looks like this:
Fewer Than Eight Electrons
Guideline: If filling the octet of the central atom
gives a - charge on central atom and a + charge on the
more electronegative outer atom:
 Don’t fill the octet of central atom.
 Most examples involve B or Be
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More than an Octet!
• There are many examples of central atoms
with more than an octet (aka expanded
valence shell or expanded octet).
• Elements of the third period and beyond
have a d subshell, so they can expand
their valence electron configurations.
– S, P, Cl (as a central atom), and other
elements in period 3 are examples
• Elements in the second period, have only s
and p subshells, so they can’t do this
More Than 8: The phosphate ion
 We can draw a Lewis structure for the phosphate ion
with an octet of 8 electrons around the central P
 We can draw a better structure with a double bond
between the P and one oxygen. Why is this better?
 When the central atom is in the 3rd row or below and
expanding its octet eliminates some formal charges,
then that is the better structure.
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Draw the Structure for IF5
Add up the valence electrons in IF5:
I has 7, F has 7
F
F
(1 x 7) + (5 x 7) = 42 electrons
F
I
F
F
I is the central atom (period 3 or greater).
Thirty-two electrons remain; first complete F octets.
The remaining 2 electrons go on I.
Covalent Bond Strength
• Bond strength is measured as how much energy it
takes to break the bond in the gas phase. The larger
the bond energy, the stronger the bond
• This is the bond enthalpy or bond energy and can
be determined experimentally.
• For example, the bond enthalpy for a Cl-Cl bond,
D(Cl-Cl), is measured to be 242 kJ/mol.
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Estimating H with Bond energies
• Bond energies can be used to estimate the enthalpy
change, H, for a reaction.
• Imagine any reaction in two steps: breaking bonds and
forming new bonds.
H = sum of bond energies for bonds broken
– sum of bond energies for bonds formed
• When H is negative, heat is released. When H is
positive, heat is absorbed.
Let’s try it! Hess’ Law approacH
• Estimate the enthalpy change for the following
reaction, using bond energies:
H
H
H
H
C
H
C
+
Cl2
H
C
C
H
H
Cl
Cl
Bonds Broken:
Bonds Formed:
1 C=C
602 kJ
1 C—C
346 kJ
1 Cl—Cl
240 kJ
2 C—Cl
654 kJ
Absorbed 842 kJ
Released 1000 kJ
H = 842 kJ – 1000 kJ
H = –158 kJ
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Notes on Bond Enthalpies
• Bond energies/enthalpies are always positive:
– Takes energy to break a bond
– Energy is always released when a bond is broken
– The bigger the bond enthalpy, the stronger the bond
• Relative bond lengths and strengths:
– Strength:
– Length:
Single bonds < double bonds < triple bonds
Single bonds > double bonds > triple bonds
Test it yourself: Length and Strength
Coming uP
• Thursday: Chapter 9
Finish Lewis Structures Worksheet in Lab
Lab Quiz 4; Green Crystals Lab due
• Tuesday: Lecture Exam 3: Chapters 6-8
Lab: Molecular Geometries Worksheet
Next Mastering Chemistry: Chap 8 due 6/27 at 11 pm
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