Chemical
Bonding
Chapter 8
1
1
Salt
•
•
•
•
NaCl
Ionic (metal+non)
Electrolyte
High MP (801ºC)
vs
Sugar
•
•
•
•
C6H12O6
Molecular (non+non)
Nonelectrolyte
Lower MP (185ºC)
2
2
Who gets with who?
• Metals + nonmetals
• IONIC
• Nonmetals + nonmetals
• MOLECULAR
• Metals + metals
• METALLIC BONDING (Alloys if a mixture.)
3
3
Ionic vs Molecular vs Metallic
•
•
•
Ionic − metals +
– Magnesium oxide
• MgO
– Potassium dichromate
• K2Cr2O7
– Nickel(II) oxide
• NiO
Molecular − elements or compounds
– Sulfur
• S8
– Bromine
• Br2
– Sucrose
• C12H22O11
Metallic − represented as monatomic
– Magnesium
• Mg
– Gold
• Au
– Copper
• Cu
4
4
Valence Electrons
• Lewis Structures
– (aka: Lewis Symbols or Dot Diagrams or Lewis
Dot Diagrams)
• Phosphorus
– Electron configuration: [Ne] 3s23p3
– Lewis Structure
– Shows only the valence electrons
– utilizes
P
5
5
Electronic Nirvana
• The octet rule - 8 valence electrons
• The reasons that atoms come to the party to
bond and form molecules is to obtain an
octet
– H only looking for a “di-tet” 2 electrons
• Using Lewis Structures to account for all the
electrons and attempt to achieve “electronic
nirvana.”
6
6
Drawing Lewis Structures
1.
2.
3.
4.
Sum the valence electrons.
Connect all atoms with single bonds
Complete the octet of all atoms.
If there are not enough electrons to go
around, use multiple bonds.
5. Put excess electrons on the central
atom even if more than an octet.
7
7
Rules of what connects to what
•
•
•
•
Sometimes the order listed may help.
Central atom is often written first.
H’s and F’s are terminal (on outside).
C’s love to hook together and do not like to be
terminal. (Avoid unshared pairs on C’s.)
• O’s do not like to bond together.
• Avoid rings of 3 or less atoms.
8
8
Lewis Structures
1. H2
6. H2S
2. O2
7. CH2O
3. N2
8. (CH3)2CO
4. CH4
9. SO2
5. NF3
10.CO2
9
9
What if there appears to be more
than one possible structure?
• First be sure you have followed the guidelines
above that would eliminate some possible
structures.
−
−
O＝C＝O
−
−
∣−
− C≡ O ∣
O
−
• Chemists have invented a tool called formal
charge that will provide information to help
eliminate some structures.
10
10
Applying Formal Charge
•
Count the number of valence electrons that “belong” to each atom.
– Unshared pairs belong entirely to the atom on which they reside.
– Shared pairs are divided evenly between the atoms on which they are shared:
half “belong” to one atom and half belong to the other atom.
•
Subtract the number of valence electrons that “belong” to each atom in the
molecule from the number of valence electrons that each atom has as an isolated
atom
(FORMULA: valence electrons in the un-bonded atom – electrons that “belong” to
the atom in the molecule)
– If the number of electrons that “belong” to the atom in the molecule is one
more than the number of valence electrons in the elemental atom, the formal
charge is -1 (two more, the formal charge is −2)
– If the number of electrons that “belong” to the atom in the molecule is one less
than the number of valence electrons in the elemental atom, the formal charge
is +1 (two less, the formal charge is +2)
– If the number of electrons that “belong” to the atom in the molecule is equal to
the number of valence electrons in the elemental atom, the formal charge is 0
−
−
O＝C＝O
−
−
0
0
0
∣−
− C≡ O ∣
O
−
-1
0
+1
11
11
Using formal charge to evaluate
the validity of structures
✓ The sum of the formal charges must equal the charge of
the species
✓ Molecules have a charge of 0
✓ A polyatomic ion’s charge should equal the formal
charge.
• Preferred structures have the lowest amount formal charge
‣ Atoms without any charge (or just the charge of the ion)
is ideal
‣ Lower charges are preferred
‣ Two charges of −1 and −1 would be better than −2 and 0
• All else being equal, any negative charge that must exist is
best placed on the most electronegative atom.
• Ultimately it may be necessary to have other information to
determine the “actual” structure.
12
12
Expanded Octet
• Elements in periods 3+ have unused dorbitals which can be accessed to put
electrons into for bonding to allow for
expanded octets.
• The maximum number of domains we will
study will be 6
– 7 and 8 do exist, but they are quite rare − and
go well beyond the scope of this course.
13
13
Lewis Structures
1. PCl5
5. SF6
2. SF4
6. BrF5
3. ClF3
7. XeF4
4. XeF2
14
14
Bonds − Length & Strength
• Single
• Overlap of one pair of electrons - Longest = weaker
• Double
• Overlap of two pair of electrons - Shorter = stronger
• Triple
• Overlap of three pair of electrons - Shortest = strongest
• more on why it is the shortest in the next chapter.
Length (A) / strength (kJ/mol)
Length (A) / strength (kJ/mol)
347
160
614
839
305
615
891
418
945
201
607
358
745
1070
204
498
15
15
Bond Polarity
• Electronegativity
– A measure of the ability of an atom (that is
bonded to another atom) to attract electrons to
itself.
– Increases up and to the right on the periodic
chart.
• To determine polarity of bond
– Subtract electronegativity of the two atoms.
– Treat each bond individually.
• Nonpolar Covalent ~(0 - 0.4)
• Polar Covalent ~(0.5 - 1.9)
• Ionic ~(2.0 +)
16
16
Determining Polarity &
Symbolizing Polarity
H
F
δ+ δH F
2.1 - 4.0 = 1.9, very polar
F is the negative end.
or
H
F
17
17
Polarity is a Continuum
•
Nonpolar covalent
•
Polar covalent
•
Ionic
18
18
Nonpolar - Polar - Ionic
• Using shape and color as a visual
representation of the three types of bonds.
• Polarity is measured as a dipole moment
measure in debye units.
0D
1.92 D
19
19
Hydrogen-halide series.
• Larger halide atoms mean smaller electronegativity.
• Difference in electronegativity decreases.
• Therefore polarity decreases and measured dipole
moment is smaller.
20
20
C6H14 (nonpolar) H2O (polar)
• C6H14 has polar bonds symmetrically arranged,
and the molecule ends up nonpolar.
• H2O has polar bonds not symmetrically
arranged, and the molecule ends up polar.
21
21
The effect of a charged rod
• The stream of polar
water coming from a
buret is bent by the
charged rod.
• Nonpolar hexane is
not bent.
22
22
Isomers
• Isomers - Molecules with the same chemical
formula, but different arrangement of atoms.
• Draw a Lewis structure for C2H6
• Draw a Lewis structure for C2H4F2
– Draw isomers and name these isomers
– Free rotation around single
• Draw a Lewis structure for C2H2F2
– Draw isomers
– Let’s name these isomers
23
23
Isomers
• Sketch and build the molecule C4H10
• Sketch and build the molecule C4H8
• Draw as many isomers as you can for this
molecule.
24
24
Resonance
• Two or more Lewis Structures in which;
– The connection and placement of the atoms is
the same
– But the arrangement of the electrons is different
25
25
Resonance in Ozone: O3
• Draw a Lewis structure for ozone.
• Re-sketch the molecule with electrons in a different
arrangement - This is called a resonance structure.
• You wouldn’t ordinarily write two different versions of this, but
chemists noticed that the two bonds in ozone were the same
length AND that they both were shorter than a single bond,
but longer than a double bond.
• Chemists think that molecule does not actually flip between
the two structures, but is an average of the two.
26
26
Resonance in N2O
• Draw two Lewis structures for N2O to
represent two different isomers.
– How do we decide which is more likely to occur
• Formal Charge
27
27
Resonance in N2O
• Draw a Lewis structure for N2O
– Put one of the N’s in the middle.
– Putting the O in the middle leads to lousy formal
charge
• Draw two other resonance structures.
• Draw an isomers of the structure above and
it’s resonance structure.
−
−
N
＝
O
＝
N
−
−
∣−
N− O≡ N ∣
−
28
28
Using Formal Charge on N2O
• Apply formal charges to the 3 resonance
structures of N2O.
0
+1
-1
-2
−
−
N ＝O＝N
−
−
-1 +2 -1
+1
+1
-1
+1
0
∣−
− O≡ N ∣
N
−
-2 +2 0
• Remember the rule:
– If there must be formal charge, the negative
formal charges should reside on the more
electronegative atoms.
29
29
Thiocyanate ions
• Look back at the six structures we drew for
the thiocyanate ions.
• Which of these are resonance structures?
(Same atom arrangement, different electron
distribution.)
• Which of these are isomers? (Different
arrangement of atoms.)
30
30
Resonance in Benzene: C6H6
• C6H6 with 30 valence electrons - make a ring structure double bonds will be necessary.
• The molecule does not flip between the two structures, but
is an average of the two.
• Bond length studies tell us that all 6 C-C bonds have the
same length - intermediate between single and double.
31
31
Representing Benzene: C6H6
• This is an alternative way of representing carbon
structures.
• At each corner the appropriate number of hydrogens are
attached.
• The circle structure on the right emphasizes the
delocalization of the electrons.
32
32
Resonance?
33
33
Exceptions to the Octet Rule
•
•
Odd number of electrons.
Sketch a Lewis Structure for
– NO, NO2
34
34
Odd Number of Electrons
Free Radicals
• Free radicals are very unstable and react quickly
with other compounds, trying to capture the
needed electron to gain stability.
• Generally, free radicals attack the nearest stable
molecule, "stealing" its electron.
• When the "attacked" molecule loses its electron, it
becomes a free radical itself, beginning a chain
reaction.
• Once the process is started, it can cascade, finally
resulting in the disruption of a living cell.
35
35
Exceptions to the Octet Rule
•
Odd number of electrons.
– NO, NO2
•
Less than an octet.
– BF3
36
36
Less Than an Octet
• Sketch a Lewis Structure for BF3
• BF3 gas is very reactive because of its
incomplete octet.
• It reacts with molecules such as water
and ammonia with their available
unshared pair of electrons.
BF3 + NH3 → H3NBF3
• In H3NBF3 the boron would now have
an octet.
37
37
Exceptions to the Octet Rule
•
Odd number of electrons.
– NO, NO2
•
Less than an octet.
– BF3
•
More than an octet.
– Expanded octet on a central atom
– only atoms that have unused “d” orbitals
– Period 3 and beyond
38
38
Lattice Energy
• The energy required to separate a mole of solid
ionic compound into its gaseous ions.
ScF3
5096
39
39
Lattice Energy
• Coulombs Law (again)
• Felectrostatic = κ Q1 Q2 / d
• Makes lattice energy
– Particularly dependent on the size of the ionic
charge
– Ionic radius is less important since ionic radii
do not vary greatly.
40
40
Early chemists describe the first dirt molecule.
41
41