Chapter 8: Bonding – General Concepts
8.1
8.2
8.3
8.4
8.5
8.7
8.9
Chemical Bond Formation
Covalent Bonding (Lewis Dot Structures)
Charge Distribution in Covalent Compounds
Resonance
Molecular Shapes (VSEPR)
Molecular Polarity
Bond Properties
Valence Electrons
Valence Electrons are in outer shell and
involved in bonding interactions
Core electrons or inner shell electrons do not
become involved with bonding interactions
and stay associated with the parent nucleus
Elements of the same group (family) have similar
chemical properties because of isoelectronic
valence shell electron configurations
Representative Elements &Valence
Shell Electron Confgurations
Group 1A:
Group 2A:
Group 3A:
Group 4A:
Group 5A:
Group 6A:
Group 7A:
Group 8A:
ns1
ns2
ns2np1
ns2np2
ns2np3
ns2np4
ns2np5
ns2np6
Representative Elements & Lewis
Dot Structures
Representative Elements & Lewis
Dot Structures
Group 1A: (ns1)
A
Group 2A: (ns2)
A
Group 3A: (ns2np1)
A
Group 4A: (ns2np2)
A
8.1 Chemical Bonds
1. Ionic Bonds - Ionic compounds between metals and
nonmetals consisting of [+] cations and [-] anions.
Group 5A: (ns2np3)
A
(ns2np4)
A
Group 7A: (ns2np5)
A
Group 8A: (ns2np6)
A
Group 6A:
2. Covalent Bonds - Classical “molecules” where
valence electrons are shared between two
nonmetals
3. Polar Covalent Bonds - Covalent bonds with ionic
character in that the electrons are not equally shared.
4. Metallic Bonds - Pure metals and alloys where
delocalized free electrons hold together the positive
nuclei.
1
Electronegativity
Electronegativity
MOST
Electronegativity- “the relative ability
of a bonding atom to attract electrons to
itself.”
LEAST
Delta Electronegativities
8.2 Covalent Bonds
Bond Type
DEN
Covalent
0.0 - 0.5
Polar Covalent
0.5 - 2.0
Ionic
4.0 - 2.0
Consider bonds to be result of atoms sharing
electron pairs
2 Types of Electron Pairs
-lone pairs
-bonding pairs
Lewis Dot Structures
-Uses dots to represent valence
electrons
-Ignores core electrons since they
are not involved in bonding
-Usually use 4 orbitals (2 e- each)
for a max of 8 e- (octet rule), note there
are important exceptions to the octet rule
Sodium = Na
Chlorine = Cl
Sodium ion = Na+
Chloride = [ Cl ]-
Exceptions to the Octet Rule
1. Atoms with odd # of electrons (free radicals)
2. Atoms which can have less than an octet
-Hydrogen can only have 2 e-Beryllium can have 4 valence e-Aluminum can have 6 valence e3. Atoms which can have expanded octets
Any atom of the 3rd or higher period has d
orbitals which can allow for more than 8 valence
electrons
2
Covalent Bonding - Diatomics
1. Hydrogen - Always has a Duet
-Is an exception to the Octet Rule
Covalent Bonding - Diatomics
3. Oxygen - has a Double Bond
or
O + O
=
or H H
2. Chlorine: (Halogens)
4. Nitrogen - has a Triple Bond
N + N
Lewis Dot Structures - Polyatomics
Represent molecular orbitals with bonding and
lone pair electrons
Accounts for all valence electrons contributed to
the molecule by each of its atoms
Resonance structures arise when there is more
than one way to arrange the electrons
Formal Charge correlates a Lewis Dot Structures
electron locations to number of valence
electrons an atom donates to a molecule
2. Determine Structural Connectivity.
N
N
or
N
N
1. Determine Total Number of Available
Electrons.
•Each atom of the representative elements donates the
number of electrons equivalent to group number
(IA – VIIA).
•Sum all available electrons, the molecule (Lewis dot
structure) must have this number of electrons
•If there is an odd number of electrons, the lone electron exists
in a nonbonding orbital or a multiple bond.
•If the compound is an ion, add one electron for each negative
charge and subtract one electron for each positive
charge
3. Draw Hypothetical structure
•First atom is central atom (unless Hydrogen, which can
never be central)
• Draw hypothetical structure with each atom surrounded by 8
electrons in pairs.
•If more than 2 types of atoms, set up in order of formula
(Acids are a notable exception).
•For each atom not surrounded by an octet of electrons (4
bonds), add lone pairs.
• Organic compounds often contain multiple carbons
connected to each other, you may need to be able
to identify functional groups.
•Never give hydrogen an octet, its 1s valence shell can only
contain 2 electrons.
•Draw a line between each bonded atom (representing
two electrons in a covalent bond)
3
4. Compare Electrons in Hypothetical
Structure to the Available Number
•Count electrons in hypothetical structure .
•Compare to the available number of electrons (from step 1).
Three Possibilities
i. Hypothetical structure has same number of
electrons: Your structure is correct
5. Can Central Atom Have Less Than an
Octet?
•Hydrogen always has a duet (one bond – 2
electrons)
•Group IIA can have a quartet (4 electrons)
•Group IIIA can have a sexter (6 electrons)
ii. Hypothetical structure has more electrons than
are available: Go to step 5.
iii. Hypothetical structure has less electrons than it
needs: Go to step 7.
6. Can Central Atom Have Less Than an
Octet?
• Place multiple bonds into structure by
replacing two loan pairs from two
adjacent atoms with a bonding pair.
• Repeat this until the structure has the same
number of electrons as determined in
step 1.
•Note, resonance structures arise if you have
more than one choice in making the
multiple bond
If not, go to step 6 and make multiple bonds
7. Can Central Atom Have Expanded
Octet?
• Atoms with d orbitals in their valence shell
(3rd period or greater) can have an
“expanded octet”, that is more than 8
valence electrons.
•Add lone pairs to these atoms until the
number of electrons equals the number
from step 1.
Class Problems
1. Ammonia
2. Ammonium Ion
3. Methane (CH4)
Group Problem
Draw Lewis Dot Structures for:
1.
2.
3.
4.
NO2
BCl3
ClF3
BrF5
4. Sulfite
4
Lewis Structure - Carbon Dioxide
# Valence e- = 4 + 2(6) = 16
O C O
This figure has 20 e-,
we only have 16,
What can we do?
Note, if we replace 2 adjacent “lone pairs” with a bonding
pair, we have reduced the total # of e- by 2 (while still
maintaining an octet), so CO2 needs 2 multiple bonds.
Resonance Structures
Resonance Structures - 2 or More Lewis
Dot Structures That Satisfy the Octet
Rule.
Consider Ozone (O3)
O1 O2 O3
O1 O2 O3
Means Resonance
Neither Structure Is Correct As All O-O Bonds Are
Identical, the True Structure Is an Average of the
Resonance Structures
Formal Charge
FC=Group # -(# of lone pair electrons
+1/2#bonding pair electrons)
Lewis Structure - Carbon Dioxide
Note, sometimes there are several
ways we can add multiple bonds
O C O
O C O
Experiments show us that
the carbon -oxygen bonds
in CO2 are identical, so
the first structure is the
best one.
O C O
8.3 Formal Charge
How do We Determine Which
Structure is Best if More Than 1
Lewis Structure is Possible?
Formal Charge:
Way of Determining If Lewis
Dot Structure Places Electrons
Near the Atom Which Donates
Those Electrons
Group Problem
Write the Lewis Dot Structure
for a Molecule of Hydrogen,
Nitrogen and Carbon
Look at the Formal Charges of each atom
-The best Lewis structure usually minimizes the
formal charge of each atom
-If structures are assigned nonzero formal charges,
the more electronegative atom should be
assigned the negative formal charge
-Polyatomic ions have atoms with non-zero formal
charges
5
Group Problem
H C
NO
FCH=1-[0+1/2(2)]=0
FCC=4-[0+1/2(8)]=0
FCN=5-[1+1/2(8)]=0
H N C
H C N
N
•These are structural isomers.
•They are isolatable and different chemical species
(resonance structures are not “real” isolatable species
H N C
FCH=1-[0+1/2(2)]=0
FCC=4-[1+1/2(8)]=-1
FCN=5-[0+1/2(8)]=1
•Use two arrows for isomers
•Use a double headed arrow for resonance structures
Which Is Best?
Consider Resonance Structures of
NO2
O=N-O
O=N-O
Shortcomings of Lewis Structures
1. Resonance Structures
O=O
O
O O=O
Incorrectly Predicts Different Bond Lengths
Which is Best?
2. Molecular Geometry
Lewis Dot Structures Can Not
Predict Molecular Geometries
O= C=O
How Do We Determine Molecular Geometries?
Two Step Process
1. Determine Valence Electron Orbital Geometry
-Note Two Types of Valence Orbitals
A. Bonding Orbitals
H O H
8.5 VSEPR
Valence Shell Electron Pair Repulsion
Theory
-Electrons in Valence Orbitals Interact to
Minimize -/- Repulsion and
Maximize +/- Attraction
B. Lone Pair Orbitals
2. Determine Molecular Geometry Based only
on Bonding Orbitals
Most Stable Geometry Maximizes Distance
Between Orbitals While Maintaining Constant
Distance From the Nucleus
6
VSEPR Geometries
Steric Number (SN)
SN= the Number of Electron Domains
Different Electronic Geometries
Result from Different Steric
Numbers
Two Electron Pairs (SN = 2)
- Linear
-180o angle
Note: Multiple Bonds are Considered
a Single Electron Domain
VSEPR Geometries
VSEPR Geometries
Three Electron Pairs (SN = 3)
4 Electron Pairs (SN = 4)
-Tetrahedral
-Trigonal Planar
-120o
-109.5o angle
angle
VSEPR Geometries
-Trigonal
Bipyramidal
5 Electron Pairs (SN=5)
VSEPR Geometries
5 Electron Pairs (SN=5)
-Trigonal
Bipyramidal
-90o axial
-90o axial
-120o equatorial
-120o equatorial
7
VSEPR Geometries
6 Electron Pairs (SN=6)
Prediction of
Molecular Geometries
1. Use VSEPR to Identify Electronic
Structure Based on 2-6 Electron Pairs
From the Lewis Dot Structure.
-Multiple Bonds Count as a
Single Electron Pair
-Octahedral
-90o
2. Visualize the Molecular Structure
From the Bonding Orbitals of the
Electronic Structure
Two Electron Pairs
Two Electron Pairs
(Linear Electronic Geometry)
(Linear Electronic Geometry)
A. 2 bonding, 0 nonbonding orbitals
A. 2 bonding, 0 nonbonding orbitals
BeH2
Three Electron Pairs
(Trigonal Planar Electronic Geometry)
BeH2
Three Electron Pairs
(Trigonal Planar Electronic Geometry)
A. 3 bonding, 0 nonbonding orbitals
A. 3 bonding, 0 nonbonding orbitals
B. 2 bonding, 1 nonbonding orbitals
BF3
8
Three Electron Pairs
Three Electron Pairs
(Trigonal Planar Electronic Geometry)
(Trigonal Planar Electronic Geometry)
A. 3 bonding, 0 nonbonding orbitals
B. 2 bonding, 1 nonbonding orbitals
BF3
CH2
Trigonal Planar Geometry
120o F-B-F bond angle
Three Electron Pairs
(Trigonal Planar Electronic Geometry)
B. 2 bonding, 1 nonbonding orbitals
Four Electron Pairs
(Tetrahedral Electronic Geometry)
A. 4 bonding, 0 nonbonding orbitals
B. 3 bonding, 1 nonbonding orbitals
C. 2 bonding, 2 nonbonding orbitals
CH2
Bent Geometry
120o H-C-H angle
Four Electron Pairs
Four Electron Pairs
(Tetrahedral Electronic Geometry)
(Tetrahedral Electronic Geometry)
A. 4 bonding, 0 nonbonding orbitals
METHANE CH4
A. 4 bonding, 0 nonbonding orbitals
METHANE CH4
-Tetrahedral Geometry
-109.5o H-C-H angle
9
Four Electron Pairs
Four Electron Pairs
(Tetrahedral Electronic Geometry)
(Tetrahedral Electronic Geometry)
B. 3 bonding, 1 nonbonding orbitals
B. 3 bonding, 1 nonbonding orbitals
Ammonia NH3
Ammonia NH3
-Trigonal Pryamidal geometry
-107o H-N-H bond angle
-Lone Pair Takes up More
Space than Bonding Pairs
Four Electron Pairs
Four Electron Pairs
(Tetrahedral Electronic Geometry)
(Tetrahedral Electronic Geometry)
c. 2 bonding, 2 nonbonding orbitals
H2 O
H2 O
Bent Geometry
O
H
c. 2 bonding, 2 nonbonding orbitals
H
Five Electron Pairs (SN=5)
(Trigonal BiPyramidal electronic geometry)
A. 5 bonding, 0 nonbonding orbitals
104.5o H-O-H bond angle
Five Electron Pairs
(Trigonal Bipyramidal Electronic Geometry)
A. 5 bonding, 0 nonbonding orbitals
PF5
B. 4 bonding, 1 nonbonding orbitals
C. 3 bonding, 2 nonbonding orbitals
D. 2 bonding, 3 nonbonding orbitals
10
Five Electron Pairs
(Trigonal Bipyramidal electronic geometry)
A. 5 bonding, 0 nonbonding orbitals
PF5
Trigonal Bipyramidal
Geometry
Five Electron Pairs
(Trigonal Bipyramidal Electronic Geometry)
C. 3 bonding, 2 nonbonding orbitals
BrF3
Five Electron Pairs
(Trigonal Bipyramidal electronic geometry)
B. 4 bonding, 1 nonbonding orbitals
SF4
See-Saw Geometry or
Distorted Tetrahedron
-Lone Pair takes
equatorial Position
Five Electron Pairs
(Trigonal BiPyramidal Electronic Geometry)
C. 3 bonding, 2 nonbonding orbitals
BrF3
T-Shaped
Five Electron Pairs
(Trigonal BiPyramidal electronic geometry)
D. 2 bonding, 3 nonbonding orbitals
XeF2
Six Electron Pairs
(Octahedral Electronic Geometry)
A. 6 bonding, 0 nonbonding orbitals
B. 5 bonding, 1 nonbonding orbitals
Linear
C. 4 bonding, 2 nonbonding orbitals
11
Six Electron Pairs
(Octahedral Electronic Geometry)
A. 6 bonding, 0 nonbonding orbitals
SF6
Six Electron Pairs
(Octahedral electronic Geometry)
A. 6 bonding, 0 nonbonding orbitals
SF6
Octahedral
Geometry
90o F-S-F bond angle
Six Electron Pairs
Six Electron Pairs
(Octahedral electronic geometry)
(Octahedral electronic geometry)
B. 5 bonding, 1 nonbonding orbitals
XeF4
IF5
F
C. 4 bonding, 2 nonbonding orbitals
F
I
F
Square Pyramidal
F
Geometry
F
Geometry of Organic Coumpounds
VSEPR Determines the Geometry of
Molecules With More Than One Central Atom
What is the geometry of each carbon
in the following compound?
HC
C
C
C
linear
8.7 Polarity of Polyatomic
Molecules
Polar Bond - Covalent Bond Between 2
Atoms of Different Electronegativities
Bond Dipole - Measure of Bond Polarity
CCH2COOH
O
H-C
Square Planar
Geometry
H
H
tetrahedral
trigonal planar
O-H
How do we Determine if a
Molecule is Polar?
bent
12
The Molecular Dipole Is the Vectoral Sum
of All the Individual Bond Dipoles
Present in a Molecule.
Does Carbon Dioxide have Polar
Bonds?
Therefor We Need to Know:
1. The Number and Types of Bonds
Is Carbon Dioxide a Polar
Molecule?
2. The Polarity of Each Type of Bond
3. The Orientation of Each Bond
Molecular Symmetry and Polarity
Is Water Polar?
Bond Polarity & Dipole Moment
Non Polar Molecules with Polar Bonds
Symmetric cancellation of Bond Dipoles
Linear Cancellation of Bond
Dipoles
What types of
geometries would
have cancellation of
bond dipoles?
13
Trigonal Planar Cancellation of
Bond Dipoles
Tetrahedral Cancellation of Bond
Dipoles
What Molecular Geometric Structures are Nonpolar
8.9: Covalent Bonds
A-central atom, B,C,D- atoms, N-nonbonding electron pairs
(l) (tp)
1. Linear(l): AB2, AB2N3
2. Trigonal Planar(tp): AB3,
3. Tetrahedral: AB4
(l) (l) (l)
4. Square Planar(sp): AB4N2, AB2C2N2
(l) (tp)
5. Trigonal Bipyramidal: AB5, AB2C3
(sp) (l)
(l) (l) (l)
6. Octahedral: AB6, AB4C2, AB2C2D2
Strength of Covalent Bonds
Bond Dissociation Energies
(kJ/mol)
Strength of Covalent Bonds
The Greater the Polarity the
Stronger the Bond of Similar Bonds
(kJ/mol)
C=O
732
O=O
498
H-H
436
C-F
450
C=O
732
O=O
498
H-C
410
C-Cl
330
H-H
436
C-F
450
C-C
C=C
C=C
350
611
835
C-Br
C-I
H-O
270
240
460
H-C
410
C-Cl
330
C-C
C=C
C=C
350
611
835
C-Br
C-I
H-O
270
240
460
- Energy Required to Split the
Atoms Apart
C-F > C-Cl > C-Br > C-I
14
Strength of Covalent Bonds
Bond Order
Multiple Bonds are Stronger Than
Single Bonds
(kJ/mol)
C=O
732
O=O
498
H-H
436
C-F
450
H-C
410
C-Cl
330
C-C
C=C
C=C
350
611
835
C-Br
C-I
H-O
270
240
460
The higher the bond order, the
stronger the bond
C C > C=C > C-C
BO:
(3)
(2)
(1)
C C > C=C > C-C
Bond Energy & Bond Length
Crystal Lattice Energies
E 
For similar types of bonds, the
greater the Energy, the shorter
the bond strength:
bond
C-C
C=C
C=C
Energy
(kJ/mol)
350
611
835
length
(nm)
14.3
12.3
11.3
Reaction Enthalpies from Bond
Dissociation Energies
DHrxn can be determine from bond enthalpies
DHrxn = S(bond dissociation energies of
bonds broken) - S(bond dissociation
energies of bonds formed)
z1 z2
d
(kJ/mol)
F-
Cl-
Br-
O-2
Li+
-1036
-853
-807
-2925
Na+
-923
-787
-747
-2695
K+
Be+2
Al+3
-821
-3505
-5215
-715
-3020
-5492
-682
-2914
-5361
-2360
-4443
-15916
Reaction Enthalpies from Bond
Dissociation Energies
Calculate DHrxn for the combustion of
methane (CH4)
CH4 + 2O2  CO2 + 2H2O
DHrxn = 4DC-H + 2DO=O -{2DC=O+4DH-O}
DHrxn = 4(410) + 2(498) –{2(732) + 4(460)}
= -668 kJ/mol
15