Ch. 9 Chemical Bonds
HW: 3,15,23,30,31,33,
35,39,43a-f,44,47a-d,50,
51,52,57
Review:
Valence electrons are the outer shell electrons of an atom.
The are the electrons that participate in chemical bonding.
Group
e- configuration
# of valence e-
1A
ns1
1
2A
ns2
2
3A
ns2np1
3
4A
ns2np2
4
5A
ns2np3
5
6A
ns2np4
6
7A
ns2np5
7
8A
ns2np6
8
Core electrons: all non-valence electrons
Review:
Chemical Bonds
• Atoms naturally form chemical bonds by linking with
other atoms to minimize energy
• Most chemical properties are determined by an atom’s
valence electrons
N-3
Al+3
1s2 2s2
2p6
1s2 2s2
2p6
Atoms seek to obtain a configuration like that
of a Noble Gas (octet) through bonding.
Electrons are attracted to the nucleus of other
atoms by electrostatic forces.
3s0
3p0
Types of Bonds
• Ionic Bond (metals & non-metals)
– e-’s are transferred from atom to atom
– Held by electrostatic attraction
• Covalent Bond (b/w non-metal atoms)
– Valence electrons are shared
• Metallic Bond (between metal atoms)
– electrons belong to no specific atom
– e-’s flow freely through the substance
TedEd: How Atom’s Bond
www.youtube.com/watch?v=NgD9yHSJ29I
Electron Dot Notation
• Dots that represent the valence electrons are
placed around the element symbol
• Maximum of 8
– Except H and He only has 2
• Must obey Hund’s Rule
– Spread out before pairing
• Not used for transition metals
Chlorine has 7
valence e-
Cl
1s2 2s2
2p5
Lewis Dot Symbols for the Representative
Elements & Noble Gases
6
The Ionic Bond
Ionic bond: the electrostatic force that holds ions
together in an ionic compound.
Li
1s22s1
+
Li+ F -
F
1s2 1s22s22p6
[He]
[Ne]
1s22s22p5
Li+ + e-
Li
+
e- +
F -
F
=
Li+ +
F -
Li+ F -
LiF
7
Electrostatic (Lattice) Energy
Lattice energy (U) is the energy required to completely separate
one mole of a solid ionic compound into gaseous ions.
Q+QE=k
r
E is the potential energy
“Coulomb’s law”
Higher ~ more stable
Lattice energy increases as
charges increase or
ion size decreases.
Q+ is the charge on the cation
Q- is the charge on the anion
r is the distance between the ions
Compound
Lattice Energy
(kJ/mol)
NaCl
MgF2
MgO
788
2957
3938
Q: +1,-1
Q: +2,-1
Q: +2,-2
LiF
LiCl
1036
853
r F- < r Cl-
The solubility rules are directly related to Lattice energy because
dissolving ionic compounds involves breaking the ions apart
Larger charge,
Stronger bond
Lattice Energy Problem
1. Predict which of the following compounds will
have the highest lattice energy?
+1 -1
NaCl
+2 -2
MgS
+2 -1
MgCl2
+1 -2
Na2S
With all period 3 atoms, the greatest ion charges hold together strongest.
2. Predict which of the following compounds will
have the highest lattice energy?
+1 -1
+1 -1
+1 -1
+1 -1
KBr
NaCl
LiF
RbI
Charges are all equal, strongest ionic bond comes from smallest ions
(Li and F atoms are in lowest period and thus smaller atoms)
The Covalent Bond
A covalent bond is a chemical bond in which two or more
electrons are shared by two atoms.
F
7e-
+
F
F F
7e-
8e- 8e-
Shared electrons are counted for both to fulfill each atom’s octet
A pair of shared electrons is represented as a line between atoms
lone pairs
Valence electrons not involved
in bonding are called Lone Pair
electrons
F
F
lone pairs
single covalent bond (2 e-)
Lewis structure of F2
Lewis structure of water
H
+
O +
H
single covalent bonds
H O H
or
H
O
H
2e- 8e- 2e-
Double bond – two atoms share two pairs of electrons
O C
O
or
O
O
C
double bonds
8e- 8e- 8e-
Triple bond – two atoms share three pairs of electrons
N N
8e- 8e-
or
Crash Course: Bonding Models and Lewis Structures:
www.youtube.com/watch?v=a8LF7JEb0IA
N
N
triple bond
12
Covalent Bonding
Atoms will normally have a number of bonds equal to the
number of valence electrons needed.
Bonds Formed
4A: 4 x C single bonds
4
3
2
1
1 x C double and 2 single bonds
1 x C triple and 1 single bond
2 x double bonds
5A: 3 x N single bonds
1 x N double and 1 single bond
1 x N triple bond
:
:
6A: 2 x O single bonds; 1 x O double bond
7A: 1 x F single bond
Lengths of Covalent Bonds
Bond Lengths
a) Dependent on atomic radius
b) Single bond > Double Bond > Triple Bond
14
Electronegativity is the ability of an atom to
attract the shared electrons in a chemical bond
toward itself.
Electron Affinity - measurable, Cl is highest
X (g) + e-
X-(g)
Electronegativity - relative, F is highest
15
The Electronegativities of Common Elements
Linus Pauling mathematically determined the measure of
attraction between an atom's nucleus and its valence electrons
by analyzing bond strength of molecules. Scale is relative
from 0 to 4.0
16
Polar covalent bond or polar bond is a
covalent bond with greater electron density
around one of the two atoms
electron rich
region
electron poor
region
H
F
Unequally shared electron density
d+
d-
H
F
d+/- indicates a
“partial charge”
Polar vs Non-Polar molecules
Polar Covalent: bonds comprised of atoms of different
electronegativity
H-F, H-Cl, H-Br, HI
Non-polar covalent: Found between bonded atoms of
similar or equal electronegativity
NCl3
O2
CH4
Crash Course: Polar & Non-Polar Molecules
www.youtube.com/watch?v=PVL24HAesnc
Classification of bonds by difference in electronegativity
Difference
Bond Type
< 0.5
Non-polar Covalent
0.5 < and < 2
Polar Covalent
2
Ionic
Increasing difference in electronegativity
Covalent
share e-
Polar Covalent
partial transfer of e-
Ionic
transfer e19
Bond Type Problem
Classify the following bonds as ionic, polar covalent, or covalent:
Strategy We follow the 2.0 rule of electronegativity difference and
look up the values
(a) the bond in HCl
(a) The electronegativity difference between H and Cl is 0.9
Therefore, the bond between H and Cl is polar covalent.
(b) The bond in KF
(b) The difference between K and F is 3.2, above the 2.0
mark; therefore, the bond between K and F is ionic.
(c) the CC bond in H3C-CH3
(d) the bond in AlP
(c) The two C atoms are identical; therefore 0,
the bond between them is Non-polar covalent.
(d) Despite being a Metal and Non-metal, the
difference is only 0.6 and is a polar covalent
Drawing Lewis Structures
• Place element with largest number of
unpaired electrons in the center
• Place other elements around the outside
• Add Lewis electron dots
– Match up the unpaired electrons
– Create Single Bonds
• If necessary, create double or triple
bonds so every atom has 8 valence
electrons.
– Hydrogen will only have 2
NF3
Lewis Structure Practice
Br2
CO2
NH3
H2CO
Bozeman Science: Lewis Diagrams
www.youtube.com/watch?v=xNYiB_2u8J4
Polyatomic ions
We add a dot (electron) for every - and remove for +
– NH4+
Net charge indicated in
brackets
– CO3-2
–
OH-
– NO2-
Many must be described by
multiple
Lewis structures (resonance)
Two possible skeletal structures of formaldehyde (CH2O)
H
C
O
H
H
C
Which is correct?
O
H
An atom’s formal charge is the difference between the number of
valence electrons in an isolated atom and the number of electrons
assigned to that atom in a Lewis structure.
Formal charge
on an atom
Valence e- in
= free atom
-
# nonbonding
electrons
-
1
2
(
total number
of bonding
electrons
)
The sum of the formal charges of the atoms in a molecule or ion must
equal the charge on the molecule or ion.
Method 2: Quicker way to find formal charge:
1. Note how many (#) electrons an atom needs to gain an octet
2. If it has the same # bonds as the #, it’s 0
If it has 1 less bond than the #, it’s -1 ; (2 less = -2)
If it has 1 more bond than the #, it’s +1 ; (2 more = +2)
1. For neutral molecules, a Lewis structure in which there are no
formal charges is preferable to one in which formal charges are
present.
2. Lewis structures with large formal charges are less plausible than
those with small formal charges.
3. Among Lewis structures having similar distributions of formal
charges, the most plausible structure is the one in which negative
formal charges are placed on the more electronegative atoms.
(b) is more likely structure
Because no formal charge
(b) is most likely
structure for N2O
Potential CH2O
structure #1
H
-1
+1
C
O
valence nonbonding 1
Formal charge =
electrons- electrons
2
Formal charge
on C
Formal charge
on O
= 4 - 2 - ½ x 6 = -1
= 6 - 2 - ½ x 6 = +1
H
(
)
bonding
electrons
Method 2
C needs 4 e- from 4 bonds, has
one less than needed so: -1
O needs 2 e- from 2 bonds, has
one more than needed so: +1
Potential CH2O
structure #2
H
H
0
C
0
O
valence nonbonding 1
Formal charge =
electrons- electrons
2
(
)
bonding
electrons
Formal charge
on C
= 4 - 0 -½ x 8 = 0
Method 2
C needs 4 e- and has 4 bonds: 0
Formal charge
on O
= 6 - 4 -½ x 4 = 0
O needs 2 e- and has 2 bonds: 0
Based on the two possible structures,
this one is favored because it has
smaller formal charges
Problem: Lewis Structure of Polyatomic ions
Write formal charges for the carbonate ion CO3-2.
C: Formal charge = 4 – 0 – 4 = 0
Oxygen in C=O: Formal charge = 6 – 4 – 2 = 0
Oxygens in C-O: Formal charge = 6 – 6 – 1 = -1
A resonance structure is one of two or more Lewis
structures for a single molecule that cannot be
represented accurately by only one Lewis structure.
Ozone
O
+
-
O
O
O
+
O
O
*Does not switch back and forth, but is a hybrid between
the two structures.
Carbonate has 3 resonance structures:
C-O bonds are 143 pm
C=O bonds are 121 pm
However, all bonds in CO3-2 are 131 pm
The -2 charge is evenly distributed across the 3 O’s
Exceptions to the Octet Rule
The Incomplete Octet
BeH2
Be – 2e2H – 2x1e4e-
H
Be
H
(Not ionic, but covalent)
B – 3e3F – 3x7e24eBF3
F
B
F
F
3 single bonds (3x2) = 6
9 lone pairs (9x2) = 18
Total = 24
30
Exceptions to the Octet Rule
Odd-Electron Molecules
N – 5eO – 6e11e-
NO
N
O
“Stable” unpaired electron:
Radical molecule, Generally
short-lived because very reactive
The Expanded Octet (central atom with n = 3+)
SF6
F
F
Can utilize 3d orbitals for
sharing more electrons
F
S
F
F
F
6 single bonds (6x2) =
12 bonding electrons
Problem: Lewis Structure with Formal Charges
Draw the Lewis structure with formal charges of:
Elements in 3p subshell (and below) can have expanded octets by utilizing
available 3d orbitals. Generally, they can form as many bonds equal to
their total number valence electrons
Phosphate
PO4-3
Sulfate
SO4-2
Perchlorate
ClO4-1
Properties of Ionic Compounds
– Generally water soluble
(polar dissolves polar)
– Dense, brittle, and hard solids
– High melting points
– Poor conductors in solid form, but good
conductors when dissolved in water or molten
Properties of Molecules
– Low melting points (dependent on size)
– Poor conductors of heat and electricity
– Fairly easy to separate in mixtures
– Exceptions to properties
• Network Covalent Substances
• 3D crystal (similar to ionic bond)
• Diamond
35
Metallic compounds
• Not limited to octet rule (d orbitals)
• Transition metals relatively stable when neutral
• Electron sea theory – low ionization energies, ecan easily move from atom to atom
– Explains the hardness of metals – difficult to separate
• Held together by delocalized electrons
– Act like “glue”
Metallic Properties
Metallic Properties– typically same as pure metals
– Conductive – because electrons can easily come travel
from one metal atom to the next metal atom.
– Luster (shine)
– Malleable