Chapter 7. Chemical Bonding I:
Basic Concepts
Chemical bond: is an attractive force that holds
2 atoms together and forms as a result of
interactions between
electrons found
in
combining atoms
• We rarely deal with isolated atoms
• Chemical bonds are broken and formed in
reactions
• Properties of substances often determined
by bonds
Bonding lowers the potential energy
How are Chemical Bonds Formed?
•
We use the electronic structure of atoms
to predict which of two types of chemical
bonds are likely:
1)Ionic (transferring electrons)
2)Covalent (sharing electrons)
•
Octet Rule: When atoms bond, they lose,
gain, or share e- to attain filled outer level
of eight e-
Ionic Bonds
-
-
Transfer of electrons from metal to nonmetal
to form ions that come close together in sold
ionic compound
- Electrostatic attractions of
closely packed, oppositely
charged ions in a regular 3-D
array
Formed when metal (that easily loses
electrons) reacts with non-metal (that easily
gain electrons).
No. e- lost by metal = No. e- gained by
nonmetal
Electron Arrangements and Ion
Charge
• Metals form
cations by losing
enough e- to form
same configuration
as previous noble
gas
• Nonmetals form
anions by gaining
enough e- to from
same configuration
as next noble gas
– Completely fill
outer s & p
orbitals
Atom
Atoms
Electron
Config
Ion
Ions
Electron
Config
Na
[Ne]3s1
Na+1
[Ne]
Mg
[Ne]3s2
Mg+2
[Ne]
Al
[Ne]3s23p1
Al+3
[Ne]
O
[He]2s22p4
O-2
[Ne]
F
[He]2s22p5
F-1
[Ne]
Properties of Ionic Compounds
• All solids at room temperature are brittle, hard, &
rigid
– Melting points greater than 300°C
• Liquid state conducts electricity, solid state does not
• If soluble in water then good electrical conductor
• Chemical formula is empirical formula - simply giving
the ratio of ions based on charge balance (no
separate molecules)
• Ions arranged in a pattern called a crystal lattice
maximizes attractions between + and – ions
Electrostatic forces and the reason ionic compounds
crack
Electrical conductance and ion mobility
Solid ionic
compound
Molten
ionic
compound
Ionic
compound
dissolved in
water
Covalent Bonds
• Atoms bond by sharing pairs of electrons
– Commonly found between nonmetal atoms
e.g. Cl2, O2, N2, H2 etc…
Example H2: overlap of 1s orbitals
2H(g) H2(g)
∆H = -432 kJ
– Most atoms form covalent bonds by sharing
enough electrons to satisfy octet rule
Bond Polarity
• Covalent bonding between unlike atoms results
in unequal sharing of the electrons
– One end has larger electron density than
other
• The result is bond polarity
– End with larger e- density gets partial “-”
charge
– End that is e- deficient gets partial “+”
charge
δ+ H
••F δ−
Dipole Moment (µ)
Electronegativity
• Relative ability of bonded atom in a molecule to attract
shared electrons
– Larger electronegativity: atom attracts more strongly
– Values 0.7 to 4.0
↑ across period (left to right) on Periodic Table
↓ down group (top to bottom) on Periodic Table
• Larger difference = more polar bond
– negative end toward more electronegative atom
Electronegativity values for
selected elements
Electronegativity and Bond Type
Difference in electronegativity (∆
∆EN)
provides a good estimate of bond type:
0 to 0.4
non-polar covalent
0.4 – 1.7
polar covalent
> 1.7
ionic
Electronegativity values for
selected elements
∆ENNaCl
= 3.0 – 0.9
= 2.1
Comparison of Ionic vs Covalent
Property
NaCl
CCl4
State
solid
liquid
Melting point (°°C)
801
-23
Molar heat of fusion
(kJ/mol)
30.2
2.5
Molar heat of vaporization
(kJ/mol)
600
30
Electrical conductivity
Good
Poor
Covalent Bonds
• Single Bond: atoms share 2 e- (1 pair)
eg., C-C
– bond order = 1
• Double Bond: atoms share 4 e- (2 pair)
eg., C=C
– bond order = 2
• Triple Bond: atoms share 6 e- (3 pair)
eg., C C
– bond order = 3
Covalent Bonds
• Bond Strength: Triple > Double > Single
– For bonds between same atoms:
C
N > C=N > C—N
• Bond Length: Single > Double > Triple
– For bonds between same atoms:
C—N > C=N > C
N
Gilbert N. LEWIS
The “Wayne Gretzky”
of chemical bonding.
American chemist: Gilbert N. Lewis (1875-1946)
Lewis Structure (Electron Dot
Symbols)
- Shows how valence electrons are arranged
among atoms in molecules & ions
- Use symbol of element to represent nucleus &
inner e- Use dots around symbol to represent valence
electrons
- Reflects idea that stability of a compound
relates to noble gas electron configuration…
- Octet Rule: Elements tend to acquire econfiguration like that of noble gases
Writing Lewis structures of
molecules
1) Count total number of valence electrons from all atoms
(add or subtract if an ion).
2) Choose the central atom.
Hint:
Central
atoms
have
the
lowest
electronegativity.
Hydrogen is always terminal while carbon is central.
3) Attach atoms together with one pair of electrons &
subtract this total.
Hint: - Line is often used for bonding electrons
4) Arrange remaining electrons in pairs:
Hint: For hydrogens, a max of 2 e- (duet) & all others
atoms have 8 e- (octet) in total around them.
Lewis Symbols
[Ne] 3s2 3p1
Al
How many valence e-s?
3
Al
[He] 2s2 2p5
F
F
IONIC BONDING
electron transfer
e.g. NaCl
The formation of ionic bonds is represented in
terms of Lewis symbols…….
Na x
+
Cl
[Na]+ [ xCl ]
Complete transference of electron to Cl- anion
21
COVALENT BONDING
electron sharing
Atoms go as far as possible toward
completing their octets by sharing electron
pairs
Consider F2
xx
x x
xF
+
xx
22
F
xx
x xF
xF
xx
Lewis Structures Practice
Write Lewis structures for the following:
Atoms:
Na,
Cl,
Ar
Ionic compounds:
NaCl,
CaF2
Covalent compounds:
CH4, C2H4, C2H2, CCl2F2, N2, CH4O,
H2NOH
Building Lewis structures of molecules
HCN as an example...
Step 1. Count the total number of valence electrons
H has 1
C has 4
N has 5
Total of 10
Step 2. Place one e- pair between each BONDED
atom
H
C
N
We have 6 e- left
All atoms must have an octet or duet
Step 3. Add electrons to terminal atoms first
The H OK it has its duet…...
24
Next...
Building
Lewis
structures
of
molecules
Step 3.
Add remaining electrons to terminal
atoms first
Add 6 electrons in pairs to give the N an octet.
H
Step 4.
C
N
Add any electrons left over to
central atom
We have none left!
Step 5. Check for an acceptable Lewis Structure
Do all atoms have an octet ???
IN THIS CASE
25
Building Lewis structures of molecules
H
C
N
No! Both C and N need an octet…..
the C and N have to share more than one pair of
ebring e- pairs from outer N atom to form
shared pairs to give C its octet!!!
Building
molecules
Step
5. CheckLewis
for anstructures
acceptable of
Lewis
Structure
bring electron pairs from outer N atom to
form shared pairs to give C its octet!!!
H
C
N
Do it again!!!!
Still no octet on C
H
H
C
N
C
N
three electron pairs between the C and N
27
Building Lewis structures of molecules
three electron pairs between the C and N
H
C
N
Lewis (electron dot) structure of HCN
There is a triple bond…..
Also written
H
C
N
Another possible structure is….
28
H
Another structure
N
C
Lets do the Lewis structure…...
Step 1. Count the total number of valence electrons
C has 4 N has 5
Step 2.
H has 1
Total of 10
Place one e- pair between each atom
H
N
C
Step 3. Place remaining electrons on terminal
atoms until their octet complete
We get……….
29
Another structure
Step 3.
Place remaining electrons on terminal
atoms until their octet complete
H
N
C
Step 4.
No electrons left.
Step 5.
Check for acceptable Lewis structure.
The N does not have an octet…...
We bring electron pairs from outer C atom
to form shared pairs to give N its octet!!!
Again we need a triple bond….
30
Lewis structure of HNC
H
N
C
three electron pairs between the C and N
this is called a triple bond…..
Also written
H
How can we choose?
N
C
H
C
The octet rule is obeyed!!….
FORMAL CHARGE………..
31
N
Formal Charge
•
Helps predict most reasonable arrangement of
atoms i.e. helps in writing Lewis structure.
•
Hypothetical charge atom would have if the
bonding electrons were shared equally.
•
Difference between no. of valence electrons
(V) on the free atom & no. assigned to atom in
molecule.
Formal charge = V – (L + ½S)
V= No. valence electrons in the free atom
L = No. of nonbonding electrons
S = No. of shared electrons
Formal Charges
Formal charge = V – (L + ½S)
• Formal charges help identify most
arrangement if more than one is possible
likely
• All formal charges of zero is best or one with
lowest number of nonzero formal charges
(closest to zero)
(a)
0
+1
-1
H
N
C
• FC (H) = 1-(0+1) = 0
• FC (N) = 5-(0+4) = +1
• FC (C) = 4-(2+3) = -1
0
0
0
(b) H
C
N
FC (H) = 1-(0+1) = 0
FC (C) = 4-(0+4) = 0
FC (N) = 5-(2+3) = 0
Using formal charges determine
which of these CO2 structures
are the most stable?
O C O
O C O
(0)
(0)
(0)
FC (O-) = 6-(6+1) = -1
FC (O=) = 6-(4+2) = 0
FC (C) = 4-(0+4) = 0
FC (C) = 4-(0+4) = 0
FC (
FC (=O) = 6-(4+2) = 0
O)
-1
= 6-(2+3) = +1
0
+1
O C O
0
0
0
O C O
(0)
(0)
(0)
Lewis Structures Practice
(a) Assign formal charges for the following
molecules:
N
S
O
O
O
O
(b) Write Lewis structures for the following
compounds:
PCl3,
NO2-,
PO43-,
CO32-
QUESTION
Oxygen difluoride is a powerful oxidizing and
fluorinating agent. Select its Lewis structure
(a)
(c)
(e)
36
F
F
O
O
F
F
None of these
(b)
(d)
F
O
F
O
F
F
Resonance Structures
• More than one Lewis structure that differs only
in position of e-s
– Lone pairs & multiple bonds in different
positions
• Actual molecule is combination (or blending) of
all resonance forms (delocalized)
– Actual structure is average of resonance
structures
•• •• ••
•• O ••
•
•• S •• O
•
..
••
•• •• ••
•
••
•• O
•
S
O
•
••
•
..
••
RESONANCE
We use a double headed arrow between the
structures..
O
O
O
O
N
N
N
O
O
O
O
O
The electrons involved are said to be
DELOCALIZED over the structure.
The blended structure is a RESONANCE HYBRID
38
QUESTION
Which of the following molecules exhibit
resonance?
1
CO2
39
2
ClO3-
3
O3
4
Cl2CO
5
F2O
Exceptions to the Lewis
Structure Rules
Some covalent molecules have central atoms that
do not have noble gas configuration because:
1) Odd No. of valence electrons (unpaired e-)
2) Central atom has fewer e- than needed for a
noble gas configuration
3) Central atom has more e- than needed for a
noble gas configuration
1. Odd No. of valence electrons
(unpaired e-)
• Molecules that contains one or more unpaired
electrons are paramagnetic or radicals.
Elements in 2nd period have only 4 orbitals in
valence shell and can’t have > 8 e- around
them.
Try to write the Lewis structure for NO
N O
2. Electron-deficient
• Few molecules contain central atoms that
don’t have filled valence shell (incomplete
octet)
• B & Be are often octet deficient with outer
atoms of hydrogen or other atoms that do
not readily form multiple bonds (BeH2)
• These are very reactive
3. Molecules with Extra
Electrons
• Elements in 3rd & higher periods (n > 2)
have more than 4 valence orbitals & can
share more than 4 pairs of electrons with
other atoms
• Expanded octet: Empty d orbitals available
& able to “expand” to 10, 12 or more e• Examples include… PCl5, SF6, ICl5
Elements in rows 3 and following can exceed the octet rule:
SF6
F
F
F
F
F
S
F
F
F
S
F
F
F
F
SF6 looks like this…….
F
F
F
How do we get Lewis Structure???
44
S
F
F
F
Chapter 8: Shapes of Molecules
• 3-D arrangement of atoms in molecule
– Structure is described by bond angle &
bond distance
• Bond angle: angle between any 2 bonds that
includes common atom (degrees)
• Bond distance: distance between nuclei of 2
bonded atoms (Å or pm)
Look at formaldehyde (H2CO)
VSEPR Theory
• Valence
shell
electron-pair
repulsion
(VSEPR) theory is used to predict molecular
geometry by examining no. of bonds &
unshared electron pairs
• KEY: Most stable arrangement is one where
valence electrons around central atom are
as far away from each other as possible
– Minimizes repulsions
Predicting Molecular Geometry
• Shape around central atom(s) can be predicted
by assuming that areas of electrons (bonding &
nonbonding) on central atom will repel each
other.
• Each bond counts as 1 area of electrons.
– single, double or triple all count as 1 area
• Each lone pair counts as 1 area of electrons
– Even though lone pairs are not attached to
other atoms, they “occupy space” around
central atom
Look at CO2
• Bonds (shared e-s) and lone
pairs of e-s are as far away
from each other as is possible
O C O
(0)
(0)
(0)
H
H C H
H
• Electron-group geometry is
linear (AX2) where A is the
central atom and X are the
terminal atoms
The electron-group
repulsions force
the groups as far
apart as possible
Tetrahedral – AX4
H
H N H
NH3 has 4 electron groups and
has a tetrahedral electrongroup geometry but it’s actual
molecular shape (molecular
geometry) is not tetrahedral
but rather trigonal-pyramidal
• VSEPR notation - AX3E where
E is the lone pair of electrons
• Angle of 109.5 °
Electron-group geometry deals with the
distribution of the electron groups
Molecular geometry deals with the molecular
shape of the molecule
Electron Group Geometry &
Molecular Geometry
The electron-domain geometry is often not the
shape of the molecule, however.
The molecular geometry is defined by the
positions of only the atoms in the molecules, not
the nonbonding pairs.
Table 8.1
Electron-Pair Geometry
•
Trigonal bipyramidal (5 regions)
– Equatorial or axial
– AX5 (trigonal bipyramidal
90°, 120°)
– AX4E (seesaw 90°, 120°),
– AX3E2 (T-shaped 90°),
– AX2E3 (linear 180°)
http://www.chemvc.com/~tim/p
hosphorus%20pentafluoride.jpg
• Octahedral (6 regions)
AX6 (octahedral, 90°)
- AX5E (square pyramidal 90°)
- AX4E2 (square planar
90°)
http://www.chemvc.com/~ti
m/boron%20pentafluoride%
20cloud%20pred.jpg
VSEPR
• All regions of high electron density are not
same
– Certain high density electron areas want more
room.
• Lone pair e-s generally spread out more than
bonding e-s, → affects bond angle → structure
looks a little different than expected.
• Order of repulsive forces:
lone pair-lone pair > lone pair-bond pair > bond
pair-bond pair
Why is the top
structure incorrect?
Electron-pair geometry vs
molecular geometry
• Electron-pair geometry: Includes all electron
pairs (this is what we just looked at)
• Molecular geometry: Includes only placement
of atoms in molecule
• Same when there are no unshared electron
pairs around central atom
• Look at methane (CH4) vs ammonia (NH3)
Molecular Geometry
• Trigonal Bipyramidal: 2 distinct positions
(i) axial: smaller
(ii) equatorial: larger
– Unshared
positions
pairs
always
occupy
equatorial
• Linear
– 2 areas of electrons around central atom,
both bonding
– Or two atom molecule is a trivial case
Predicting a VSEPR Structure
1. Draw Lewis structure.
2. Count number of regions of high e- density
(unshared pairs and bonds) around central
atom.
3. Identify electron-pair geometry.
4. If more than one arrangement is possible,
choose one that minimizes unshared pair
repulsions.
Multiple Covalent Bonds
Predict the electron geometry and molecular
geometry of SO2
O
S
O
O
S
O
Molecules With More Than One
Central Atom
What is the electron-group geometry of methyl
isocyanate, CH3NCO?
Ans: Draw the best fit Lewis structure
Valence e-s: C = 8
N=5
O=6
H=3
Total = 22
H
H C N C O
H
H
H C N C O
H
o
180
N
H
C
120o
C
H
O
H
109o
Practice
Predict electron-pair & molecular
geometries for following:
SiCl4, H3O+, SF4, XeF4, CH3OH,
NH2CH2COOH