Chapter 8
“Covalent
Bonding”
Covalent Bonds
• covalent
• co- (Latin - “together”)
• valere - “to be strong”
• 2 e- shared - strength holding 2
atoms together
• Smallest particle - “molecule”
Molecules
 Some elements in nature are
molecules:
 neutral group of atoms covalently
bonded
 Ex. - air contains O molecules, 2 O atoms
joined covalently
 Called “diatomic molecule” (O2)
How does H2 form?
(diatomic hydrogen molecule)
• The nuclei repel each other, (both
have + charge)
+
+
+
+
How does H2 form?
• nuclei attraction to e-’s stronger
than repulsion of nuclei
• e-’s shared
• covalent bond
• Only NONMETALS!
+
+
Covalent bonds
• Nonmetals
• don’t lose e-’s
• still want NGC
• share valence e- w/ each other = covalent
bonding
• both atoms count shared e- for NGC
Covalent bonding w/ Fluorine atoms
Covalent bonding
Fluorine has 7 valence e A second atom also has seven
 By sharing electrons…
 …both end with full orbitals

F F
8 Valence
electrons
Covalent bonding
Fluorine has 7 valence e A second atom also has 7
 By sharing electrons…
 …both full orbitals

8 Valence
electrons
F F
single covalent bond between 2 H atoms
Molecular Formulas
• molecular compounds - Compounds
bonded covalently
• Molecular compounds have:
• lower melting/boiling pts
• Generally stronger bond than ionic
• gases or liquids @ RT
• molecular formula:
• Shows # atoms of each element in
molecule
Reminder from Ch. 7
• No “molecule” of sodium chloride
• Ionic cmpds exist as collection of + & charged ions arranged in repeating 3D
patterns.
• Formula unit
Molecular Formulas
• water H2O
• Subscript “2” means 2 atoms of H
• subscript 1 omitted
• Molecular formulas do not give info
about structure (arrangement of atoms)
- Page 215
These are some of the
different ways to represent
ammonia:
1. The molecular
formula shows
how many atoms
of each element
are present
2. The structural
formula ALSO
shows the
arrangement of
these atoms!
3. The ball and stick model is
BEST, because it shows 3D
arrangement.
Section 8.2
The Nature of Covalent Bonding
A Single Covalent Bond is...
• sharing 2 valence e• Only nonmetals and H
• Different from ionic bond b/c they
actually form molecules.
• Two specific atoms joined
• In an ionic solid, you can’t tell which
atom e-’s moved from or to
How to show the formation…
• It’s like a jigsaw puzzle.
• You put the pieces together to end up with the right formula.
• C ….. special ex. - can it really share 4 electrons: 1s 2s 2p ?
2
2p
1s
2s
C
•Yes, due to e- promotion!
2
2
How to show the formation…
• It’s like a jigsaw puzzle.
• You put the pieces together to end up with the right formula.
• Carbon is a special example - can it really share 4 electrons:
1s22s22p2?
2p
1s
2s
Water
Another example: water formed w/ covalent bonds, by
using e- dot structure
H
O
Each H has 1 valence e- Each H wants 1
O has 6 valence e- wants 2 more
Sharing completes each
octet
Single bonds
Water
• Put the pieces together
• The first H is happy
• The O still needs 1 more
HO
Water
• a 2nd H attaches
• Every atom has full energy levels
HO
H
Note the two
“unshared” pairs
of electrons
This is what
happens when
you miss
Chemistry
class
Do you
know this
kid?
How many bonds do I make?
#bonds = # valence e- needed # valence e- present / 2
Check out my webpage for a detailed list
of rules for electron dot/structural
formula rules for covalent bonds
Examples:
1. Conceptual Problem
8.1 on page 220
We’ll do #7 & 8
Double and Triple Covalent Bonds
• Sometimes atoms share more than one
pair of valence e-’s
• double bond: atoms share 2 pairs of e-’s
(4 total)
• triple bond: atoms share 3 pairs of e-’s
(6 total)
• Table 8.1, p.222 - Know these 7
elements as diatomic:
Br2 I2 N2 Cl2 H2 O2 F2
Dot diagram for Carbon dioxide
C
• CO2 - Carbon is central
atom ( more metallic )
• C - 4 valence e• Wants 4 more
• O- 6 valence e-
O
• Wants 2 more
The chemistry of CO2 6:44
Carbon dioxide
• Attaching 1 oxygen leaves oxygen 1 eshort, and carbon 3 short
CO
Carbon dioxide
2nd O
 both O 1 e- short
 C 2 e- short
 Attaching
OC O
Carbon dioxide
 only
solution  share more
O CO
Carbon dioxide
 The
only solution is to share more
O CO
Carbon dioxide
 The
only solution is to share more
O CO
Carbon dioxide
 The
only solution is to share more
O C O
Carbon dioxide
 The
only solution is to share more
O C O
Carbon dioxide
 The
only solution is to share more
O C O
Carbon dioxide
 The
only solution is to share more
 Requires 2 double bonds
 Each atom counts all e- in bond
O C O
Carbon dioxide
The only solution is to share more
 Requires two double bonds
 Each atom counts all e- in the bond

8 valence
electrons
O C O
Carbon dioxide
The only solution is to share more
 Requires two double bonds
 Each atom counts all e- in the bond

8 valence
electrons
O C O
Carbon dioxide
The only solution is to share more
 Requires two double bonds
 Each atom counts all e- in the bond

8 valence
electrons
O C O
How covalent bonds form - Mark Rosengarden
How to draw them?
 Use the handout guidelines:
1) Add up all valence e2) Count total e- needed to make all atoms
happy (stable)
3) Subtract; Divide by 2 (tells you how many
bonds to draw)
4) Central atom (least electronegative)
5) Start w/ most electronegative atom, fill in
remaining valence e- to fill atoms up
Examples
N
H
• NH3, ( ammonia )
• N – central atom; 5 valence e(wants 8)
• H - has 1 (x3) valence
electrons, wants 2 (x3)
• NH3 has 5+3 = 8
• NH3 wants 8+6 = 14
• (14-8)/2= 3 bonds
• 4 atoms with 3 bonds
Examples
• Draw in the bonds; start with singles
• All 8 e- accounted for
• Everything full – DONE!
H
H NH
Example: HCN
• HCN: C is central atom
• N - has 5 valence electrons, wants 8
• C - has 4 valence electrons, wants 8
• H - has 1 valence electron, wants 2
• HCN has 5+4+1 = 10
• HCN wants 8+8+2 = 18
• (18-10)/2= 4 bonds
• 3 atoms with 4 bonds – this will require
multiple bonds - not to H however
HCN
• Put single bond between each atom
• Need to add 2 more bonds
• Must go between C and N (Hydrogen is full)
HC N
HCN
Put in single bonds
 Needs 2 more bonds
 Must go between C and N, not the H
 Uses 8 electrons – need 2 more to
equal the 10 it has

HC N
HCN
Put in single bonds
 Need 2 more bonds
 Must go between C and N
 Uses 8 electrons - 2 more to add
 Must go on the N to fill its octet

HC N
Another way of indicating
bonds
• Often use a line to indicate a bond
• structural formula
• Each line = 2 valence e-
HOH H O H
=
Other Structural Examples
H C N
H
C O
H
A Coordinate Covalent Bond...
• When one atom donates both
electrons in a covalent bond.
• Carbon monoxide (CO) is a good
example:
Both the carbon
and oxygen give
another single
electron to share
CO
Coordinate Covalent Bond
 When
one atom donates both e i.e. Carbon monoxide (CO)
This carbon
electron
moves to
make a pair
with the other
single.
C O
Oxygen
gives both of
these
electrons,
since it has
no more
singles to
share.
Coordinate Covalent Bond
 When
one atom donates both e Carbon monoxide (CO)
The
coordinate
covalent bond
is shown with
an arrow as:
C
O
C O
Coordinate covalent bond
• Most polyatomic cations & anions
contain covalent & coordinate
covalent bonds
• Table 8.2, p.224
• Sample Problem 8.2, p.225
• Ammonium ion (NH4+) can be
shown as another example
Bond Dissociation Energies...
• Total energy required to break bond
btwn 2 covalently bonded atoms
• High dissociation energy usually
means compound relatively
unreactive, b/c high energy
needed to break bond
Resonance is...
• When more than one valid dot diagram
is possible.
• Consider the two ways to draw ozone (O3)
• Which one is it? Does it go back and
forth?
• It’s hybrid of both, shown by doubleheaded arrow
• found in double-bond structures!
Resonance in Ozone
Note the different location of the double bond
Neither single structure is correct,
actually a hybrid of the two. To show
it, draw all possible structures, and
join them with a double-headed arrow.
Resonance
Occurs when more than one valid Lewis structure
can be written for particular molecule (due to
position of double bond)
•resonance structures of carbonate ion (used
in production of carbonated beverages).
•The actual structure is an avg (or hybrid) of
these structures.
Polyatomic ions – note the different
positions of the double bond.
Resonance
in a
carbonate
ion (CO32-):
Resonance
in an
acetate ion
(C2H3O21-):
The 3 Exceptions to Octet rule
• For some molecules, it is impossible to
satisfy the octet rule
#1. usually when there is an odd
number of valence electrons
• NO2 has 17 valence electrons,
because the N has 5, and each O
contributes 6. Note “N” page 228
• It is impossible to satisfy octet rule, yet
the stable molecule does exist
Exceptions to Octet rule
•Another exception: Boron
•Page 228 shows boron trifluoride, and
note that one of the fluorides might be
able to make a coordinate covalent bond
to fulfill the boron
•#2 -But fluorine has a high
electronegativity (it is greedy), so this
coordinate bond does not form
•#3 -Top page 229 examples exist because
they are in period 3 or below
Section 8.3
Bonding Theories
• OBJECTIVES:
•Describe the relationship
between atomic and
molecular orbitals.
Section 8.3
Bonding Theories
• OBJECTIVES:
•Describe how VSEPR theory
helps predict the shapes of
molecules.
Molecular Orbitals are...
• The model for covalent bonding
assumes orbitals are those of individual
atoms = atomic orbital
• Orbitals that apply to overall molecule,
due to atomic orbital overlap are
molecular orbitals
• bonding orbital is molecular orbital
occupied by 2 e-’s of covalent bond
Molecular Orbitals - definitions
• Sigma bond- 2 atomic orbitals combine
to form molecular orbital symmetrical
along axis connecting nuclei
• Pi bond- bonding e-’s likely above and
below bond axis (weaker than sigma)
• Note pictures - next slide
Hybridization video – 1:36
- Pages 230 and 231
Sigma bond is symmetrical along
axis between 2 nuclei.
Pi bond is
above and
below bond
axis weaker
than sigma
Attractive & repulsive forces in H2
bond
• + nuclei repel
• e-’s repel
• Nuclei and e-’s attract
• Attractive forces stronger than repulsion
• As nuclei distances decrease, PE decreases
• If distance decreases more, PE increases
b/c increased repulsion
• Bond forms w/ bond length = interatomic
distance (PE minimum)
VSEPR: stands for...
• Valence Shell Electron Pair Repulsion
• Predicts 3D shape of molecules
• The name tells you the theory:
• Valence shell = outside e-’s
• Electron Pair repulsion = e- pairs try to get
as far away as possible from each other.
• determines angles of bonds.
VSEPR
• Based on # of pairs of ve-’s, bonded &
unbonded.
• Unbonded pair called lone pair.
• CH4 - draw structural formula
• Has 4 + 4(1) = 8
• wants 8 + 4(2) = 16
• (16-8)/2 = 4 bonds
VSEPR for methane (a gas):
• Single bonds fill
all atoms.
• 4 pairs of e-’s
pushing away
• The furthest
they can get
away is 109.5°
4 atoms bonded
• Basic shape tetrahedral
• pyramid w/
triangular base
• Same shape for
everything with 4
pairs
H
H
C
H
109.5º
H
Other angles, pages 232 - 233
• Ammonia (NH3) = 107o
• Water (H2O) = 105o
• Carbon dioxide (CO2) =
o
180
VSEPR models
VSEPR theory video 4:52
- Page 232
Methane has
an angle of
109.5o, called
tetrahedral
Ammonia has
an angle of
107o, called
pyramidal
Note the unshared pair that is repulsion for other electrons.
VSEPR song – 4:33
Hybrid Orbitals
• Provides info for molecular bonding & shape
Hybridization w/ single bonds
• Orbitals combine
• C outer e- configuration 2s2 2p2 but one 2s epromoted to 2p
•
•
•
•
One 2s e- and 3 2p e-’s
Allows bond in methane (CH4)
All bonds same due to orbital hybridization
Mix to form 4 sp3 hybrid orbitals
Hybridization w/ double bonds
• Ethene C2H4 – 1 C-C double bond and 4 C-H single bonds
• sp2 hybrid orbitals form from one 2s and two 2p atomic
orbitals of C
Section 8.4
Polar Bonds and Molecules
• OBJECTIVES:
•Describe how
electronegativity values
determine the distribution
of charge in a polar
molecule.
Section 8.4
Polar Bonds and Molecules
• OBJECTIVES:
•Describe what happens to
polar molecules when they
are placed between
oppositely charged metal
plates.
Section 8.4
Polar Bonds and Molecules
• OBJECTIVES:
•Evaluate the strength of
intermolecular attractions
compared with the strength
of ionic and covalent bonds.
Section 8.4
Polar Bonds and Molecules
• OBJECTIVES:
•Identify the reason why
network solids have high
melting points.
Bond Polarity
• do covalent bonds always share
equally?
• e-’s pulled - tug-of-war, btwn nuclei
• In equal sharing (such as diatomic molecules), the
bond is called nonpolar covalent bond
Bond Polarity
• When 2 different atoms bond covalently,
unequal sharing
• more electronegative atom stronger
attraction
• slightly negative charge
•polar bond
Lower case delta
Electronegativity?
The ability of an atom in a
molecule to attract shared
electrons to itself.
Linus Pauling
1901 - 1994
Bond Polarity
• Refer to periodic table w/ EN
• Consider HCl
H = electronegativity of 2.1
Cl = electronegativity of 3.0
• Polar bond
• Cl: slight - charge
• H: slight + charge
Bond Polarity
• Partial charges, much less than 1+ or
1- in ionic bond
d+
d-
H Cl
Partial charges
+
d and
d
Bond Polarity
• Can also be:
H Cl
• arrow points to more EN atom
KNOW Table 8.3, p.238 shows how
electronegativity indicates bond type
Polar molecules
• Sample Problem 8.3, p.239
• polar bond tends to make entire
molecule “polar”
• areas of “difference”
• HCl has polar bonds, thus polar molecule
• molecule w/ 2 poles called dipole, like
HCl
Polar molecules
• effect of polar bonds on polarity of entire
molecule depends on molecule shape
• CO2 has 2 polar bonds
• linear
• nonpolar molecule
Polar molecules
• effect of polar bonds on molecule polarity
depends on shape
• water has 2 polar bonds - bent shape
• highly electronegative O pulls e- away from H
• very polar!
polar bond of water molecule
Chemistry of water 4:46
bonding animations
Polar and non-polar covalent bonds - Mark Rosengarten
Attractions between molecules p.
240
• makes solid & liquid molecular cmpds possible
• weakest called van der Waal’s forces - there
are two kinds:
#1. Dispersion forces: weakest of all, caused by
motion of e- - increases as # e- increases
(momentarily more on side of molecule closest
to neighboring molecule, neighboring
molecule’s e-’s move to opp side)
halogens start as gases; bromine (l); iodine (s)
– all in Group 7A
#2. Dipole interactions
• Occurs when polar
molecules attracted to each
other
• 2. Dipole interaction
happens in water Figure
8.25, page 240
• +region of one molecule
attracts -region of another
molecule
#2. Dipole interactions
• Occur when polar molecules attracted to each
other
• Slightly stronger than dispersion forces
• Opposites attract, but not completely hooked
like ionic solids
+
d
d
H F
+
d
d
H F
#2. Dipole Interactions
+
d
d+
d-
d
#3. Hydrogen bonding
• …is the attractive force caused by
hydrogen bonded to N, O, F, or Cl
• N, O, F, and Cl very electronegative, so
this is very strong dipole
• And, the hydrogen shares with lone pair in
molecule next to it
• This is strongest of the intermolecular
forces
Order of Intermolecular attraction strengths
1) Dispersion forces are weakest
2) Little stronger are dipole
interactions
3) Strongest is H bonding
4) All are weaker than ionic bonds
#3. Hydrogen bonding defined:
• When H atom is: a) covalently bonded to a
highly EN atom, AND b) is also weakly
bonded to unshared e- pair of nearby highly
EN atom
• The H is left very e- deficient (only had 1 to
start with!) - it shares with something
nearby
• H is ONLY element with no shielding for its
nucleus when involved in covalent bond!
Hydrogen Bonding
(Shown in water)
d+ dH O
+
Hd
This H is bonded
covalently to: 1) the highly
negative O, and 2) a
nearby unshared pair.
H bonding allows H2O to be a liquid at
room temp
H O
H
Attractions and properties
• Why are some chemicals gases,
some liquids, some solids?
• Depends on type of bonding!
• Table 8.4, page 244
• Network solids – solids where all
atoms covalently bonded to each
other
Attractions and properties
• Figure 8.28, page 243
• Network solids melt at very high
temps, or not at all (decomposes)
• Diamond does not really melt, but
vaporizes to a gas at 3500 oC
• SiC, used in grinding, has melting pt
of 2700 oC
Covalent Network Compounds
Some covalently bonded substances DO
NOT form discrete molecules.
Diamond, a network of
covalently bonded carbon
atoms
Graphite, a network of
covalently bonded carbon
atoms
Ionic/Covalent Bond Song