BONDING
Unit 4
OVERVIEW



Valence electrons
Ions (cations/anions)
Types of bonds







Octet Rule, covalent
compounds, exceptions,
ions, resonance, formal
charge, isomers
Bond order, energy, length
Ionic Lewis Structures
VSEPR Theory


Bond strength,
hardness/texture, melting
point/boiling point,
conductivity
Metallic Bonds
Lattice Energy
Electronegativity & Bonds
Lewis Dot Structures


Ionic, covalent
Properties of bonds


Molecular Geometry
Intermolecular Forces
Hydrogen
Ion-dipole
Dipole-dipole
 London dispersion




Valence Bond Theory


Hybridization
Sigma/pi bonds
HOW ARE COMPOUNDS FORMED?


Electrons are lost, gained, or shared.
Valence electrons: the electrons available to be
lost, gained, or shared
Number of valence electrons is electrons in s and p
orbitals of highest energy level (outermost shell)
 Main groups on table correspond to number of
valence electrons.

IONS

Atom(s) that has a negative (-) or positive (+) charge

Cation: positively charged ion
Na

+
Na
Anion: negatively charged ion
e- + Cl
Cl-
+
e
CATIONS
Forms by losing electrons
 Atoms with less than 4 valence electrons
generally lose electrons

ANIONS
Forms by gaining electrons
 Atoms with more than 4 valence electrons
generally gain electrons

IONIC BOND


Electrostatic attraction between cation and anion (opposite
charges attract)
 For example, an uncharged chlorine atom can pull one
electron from an uncharged sodium atom, yielding
Cl−and Na+.
 Cation becomes smaller and anion becomes larger
Typically between a metal and nonmetal
IONIC BONDS

Ionic compounds do not exist as single molecules

Form crystal lattice
Solid crystal at ordinary
temperatures
 Organized in a characteristic
pattern of alternating positive
and negative ions


All salts are ionic compounds
and form crystals
COVALENT BOND

Bond involves sharing of electrons

Forms between atoms of two nonmetallic elements

Forms molecules

Nonpolar covalent bond: electrons are shared equally
resulting in an even distribution of the negative charge
(no partial negative charge)

Polar covalent bond: one atom in the bond attracts
electrons more than the other atom making the electron
negative charge shift to that atom giving it a partial
negative charge
BOND
TYPES
BOND STRENGTH
IONIC


Strong bonds
Crystal lattice
 Repeating symmetrical
pattern
COVALENT



Weaker bonds
Covalent Network
Molecules
 Strong forces within molecule
but weak between molecules
MELTING & BOILING POINT
IONIC


High melting/boiling
points
 Sturdy crystal lattice
Low volatility
COVALENT



Low melting/boiling points
 Weak forces between
molecules
Most are gases
High volatility
TEXTURE & HARDNESS
IONIC


Brittle
 Sturdy but collapses
easily if disrupted
Hard structure
COVALENT


Soft compounds
Most are gases
CONDUCTIVITY – IONIC
IONIC



COVALENT
Solids – not conductors
 Not good conductors
 Ions can’t move
Molten state – conductors
Solution – conductors (ions separate)
METALLIC BOND

Results from attraction
between metal cations and
the surrounding “sea of
electrons”


Vacant p and d orbitals in
metal's outer energy levels
overlap, and allow outer
electrons to move freely
throughout the metal
Valence electrons do not
belong to any one atom
METALLIC BONDS

Very strong bonds


Due to “sea of electrons”
Highest melting/boiling points

Electrons result in strong forces holding together
Low volatility
 Very hard

METALLIC BONDS

Malleable/ductile


Great conductors (heat and electricity)


When struck, one plane of metal atoms can slide past
another without breaking
Freedom of electrons carries current
Shiny and have luster

Electrons absorb light,
get excited, fall, re-radiate
the light
LATTICE ENERGY
The energy required to separate ions of an ionic
solid
 Magnitude of lattice energy depends on

Charges of ions
 Sizes of ions
 Arrangement in ions
in the solid

For given arrangement of
ions, lattice energy
increases as the charges on
the ions increase and as
their radii decrease
Compound
LiCl
NaCl
KCl
NaBr
Na2O
Na2S
MgCl2
MgO
Lattice energy
kJ/mol
834
769
701
732
2481
2192
2326
3795
ELECTRONEGATIVITY
Atom’s ability to attract electrons
 Difference in electronegativity values between
two elements can determine type of bond

Ionic Bonding = Over 1.7
 Polar Covalent = 1.7 - 0.4
 Non-polar Covalent = 0-0.3


Examples
KCl = 2.2 (ionic)
 CuS = 0.6 (polar covalent)
 O2 = 0.0 (nonpolar covalent)

LEWIS DOT NOTATIONS

Valence electrons are represented by dots drawn
around the symbol of an element
1 valence e-
2 valence e-
3 valence e-
4 valence e-
X
X
X
X
5 valence e-
6 valence e-
7 valence e-
8 valence e-
X
X
X
X
OCTET RULE

Compounds tend to form so that each atom, by
gaining, losing, or sharing electrons, has 8
electrons in its outer shell.

Hydrogen is exception and only needs 2 electrons
for a complete shell
Electrons!
LEWIS DOT STRUCTURES

Show electron distribution in compounds
Lone pair: nonbonding pair of valence electrons
 Bond pair: valence electrons shared between two
elements


Shared electrons pairs represented by two dots (:)
or by a single line ( - )
••
•
H • Cl ••
••
shared or
bond pair
lone pair
STEPS FOR BUILDING A DOT STRUCTURE
Ammonia, NH3
1. Add up the total number of valence electrons that can be
used.
H = 1 and N = 5
Total = 1 + 1 + 1 + 5 = 8 electrons
2. Decide on the central atom;
If carbon it is always central atom
If no carbon, choose the least electronegative atom
Never use hydrogen!
Therefore, N is central on this one
Building a Dot Structure
3. Draw the basic skeletal
structure of the molecule.
H N H
H
4. Draw in the electrons around the
central atom. Then place the remaining
electrons to form complete octets and
lone pairs as necessary.
N has 3 BOND PAIRS and 1 LONE PAIR.
.. .
.
H . .N. . H
H
Building a Dot Structure
5.
Check to make sure there are 8 electrons
around each atom except H.
..
.
.
.
H
H .N
.H.
6. Count electrons! Make sure that the number of
electrons in your structure equals the total number from
the beginning.
Total Electrons = 8
..
.
.
.
H .N
.. H
H
MULTIPLE BONDS (DOUBLE)
H
C
H
H
H
OR
C
H
H
C
H
Example: ethene
Two pairs of shared electrons
C
H
MULTIPLE BONDS (TRIPLE)
H
C
C
H
C
H
OR
H
C
Example: ethyne
Three pairs of shared electrons
EXCEPTIONS TO THE OCTET RULE
Element is 3rd period or higher, is the central
atom, AND is bonded to electronegative atoms
(such as O, F, Cl, Br) may have more than 8
electrons


Electrons use empty valence d orbitals

Be is stable with 4 electrons

B is stable with 6 electrons
LEWIS DOT STRUCTURES FOR IONS
 Add


or remove electrons based on charge of ion
If the ion has a negative (-) charge add electrons to
the Lewis structure.
If the ion has a positive (+) charge, then subtract
electrons from the Lewis structure.
-2
Try CO3-2
O
C
O
O
SPECIES WITH AN ODD
TOTAL NUMBER OF ELECTRONS
A very few species exist where the total number
of valence electrons is an odd number
 This must mean that there is an unpaired
electron which is usually very reactive.


Radical
Species that has one or more unpaired electrons
 They are believed to play significant roles in aging
and cancer.

Example – NO

It has a total of 11 valence electrons:
6 from oxygen and 5 from nitrogen.
:

.
:N::O:
Draw O3
O O
O O
O
O
O
O O
O
O
O
O
O
O
Which are the same and which are different?
RESONANCE STRUCTURES
Occurs when more than one valid Lewis
structure can be written for a particular
molecule.
 The resonance structures are the same except for
the placement of the electrons (meaning the
bonds).

RESONANCE BONDS
Resonance bonds are shorter and stronger than
single bonds.
H
H
H
H
H
H
H
H
H
H
H
H
Resonance bonds are longer and weaker than
double bonds.
FORMAL CHARGES

A bookkeeping system for electrons.
Does not give actual charge of atoms
 Helps decide which Lewis structure is more preferred



Used to show the approximate distribution of electron
density in a molecule or polyatomic ion
Assign an atom electrons:


Each atom gets half of the bonding electrons it has
(single, double, and triple bonds)
Each atom gets all unshared (nonbonding) electrons that
are found on it
FORMAL CHARGES

To solve for formal charge:
(Total valence electrons) – (electrons assigned to atom)

To decide which structure is preferred:

Choose the Lewis structure in which the atoms bear
formal charges closest to zero

Choose the Lewis structure in which any negative
charges reside on the more electronegative atoms
FORMAL CHARGES

Example: CO2
O=C=O
For each oxygen
4 electrons from unshared electrons
2 from the bonds (1/2 of 4)
6 total
Formal charge = 6 - 6 = 0
For carbon
0 unshared electrons
4 from the bonds (1/2 of 8)
4 total
Formal charge = 4 - 4 = 0
ISOMERS
Isomers – compounds whose molecules have the
same overall molecules but different structures
H H
H C C O H
H
H
H C O C H
H H
H
H
H H
H
H
H C C O H
H H
H C O C H
H
H
BOND ORDER

Refers to the average number of bonds that an
atom makes in all of its bonds to other atoms
Draw Lewis structures and determine bonds
 Single bond counts as 1 bond, double counts as 2, and
triple counts as 3

Bond Order =
Examples:
CH3Cl = 4/4 = 1
CO3-2 = 4/3 = 4/3
number of bonds
number of atoms bonded
CO2 = 4/2 = 2
ClO4-1 = 4/4 = 1
BOND ENERGIES

Bond energies and lengths differ between single,
double, and triple bonds
 Bonds between elements become shorter and
stronger as multiplicity increases.
 The greater the bond energy, the shorter the bond
length
Bond
type
C
C
C
C
C
C
Bond
order
Length
pm
1
2
3
154
134
120
Bond energy
kJ/mol
347
615
812
BOND LENGTH AND ENERGY
Bond
C-C
C=C
Length (pm)
154
134
Energy (kJ/mol)
346
612
CC
C-N
C=N
120
147
132
835
305
615
CN
C-O
C=O
116
143
120
113
887
358
799
1072
145
125
110
180
418
942
CO
N-N
N=N
NN
IONIC COMPOUNDS & LEWIS STRUCTURES
Na +
Cl
Na + Cl
+
The electron from Na is given to the Cl. Now both
satisfy the octet rule.
Below, two electrons are given to S, one from each K.
VSEPR model
Valence Shell Electron Pair Repulsion theory
Molecule adopts the shape
that minimizes the electron
pair repulsions.
Most important factor in determining geometry is
relative repulsion between electron pairs.
Electrons arranged so that pairs are as far apart
from each other as possible.
Occurs in 3-dimensional space
MOLECULAR GEOMETRY

Molecules have specific shapes.

Determined by the number of electron pairs around
the central species
• Bonded and unshared pairs count (unshared pairs
take up slightly more space)
• Multiple bonds are treated as a single bond for
geometry.
•Geometry affects factors like polarity and
solubility.
1 ELECTRON DOMAIN

1 atom bonded to another atom
Electron
Domains
Basic
Geometry
0 lone pair 1 lone pair
1
Linear
180°
2 lone
pairs
3 lone
pairs
4 lone
pairs
2 ELECTRON DOMAINS

2 atoms, or lone electron pairs, or a combination
of the two, bonded to a central atom.
Electron
Domains
2
Basic
Geometry
0 lone pair
1 lone pair
Linear
180°
Linear
180°
2 lone pairs
3 lone pairs
3 ELECTRON DOMAINS

3 atoms, or lone electron pairs, or a combination
of the two, bonded to a central atom.
Electron
Domains
3
Basic
Geometry
0 lone pair
1 lone pair
trigonal planar
120°
bent / angular
<120°
2 lone pairs
3 lone pairs
4 ELECTRON DOMAINS

4 atoms, or lone electron pairs, or a combination
of the two, bonded to a central atom.
Electron
Domains
Basic
Geometry
0 lone pair
1 lone pair
Tetrahedral
109.5°
Pyramidal
<109.5°
2 lone pairs
4
bent / angular
<109.5°
5 ELECTRON DOMAINS

5 atoms, or lone electron pairs, or a combination
of the two, bonded to a central atom.
Electron
Domains
5
Basic
Geometry
0 lone pair
1 lone pair
trigonal
bipyramidal
180°, 120°, 90°
seesaw/
sawhorse
2 lone pairs
t-shape
3 lone pairs
linear
6 ELECTRON DOMAINS

6 atoms, or lone electron pairs, or a combination
of the two, bonded to a central atom.
Electron
Domains
Basic
Geometry
0 lone pair
1 lone pair
2 lone pairs
Octahedral
90°, 180°
square
pyramid
square planar
6
3 lone pairs
INTERMOLECULAR FORCES

Forces of attraction between molecules

Forces within molecules are intramolecular
Why would boiling point be a good
indicator of intermolecular force strength?
HYDROGEN BONDING

H bonded to N, O, F


Example: Water
Strongest intermolecular force
Large electronegativity difference
 Size of H atom allows it to get close to unshared pair
of electrons in adjacent molecule

POLARITY

A molecule, such as HF, that has a center of
positive charge and a center of negative charge is
said to be polar, or to have a dipole moment.
H F
+
Slight positive side
Smaller electronegativity
Slight negative
Larger electronegativity
DIPOLES

Dipole – created by equal but opposite charges
that are separated by a short distance
Molecules with dipoles are polar because of uneven
charge distribution
 Direction is from its positive to negative pole
 Molecules can have multiple dipoles
 Dipoles can cancel each other out if in equal and
opposite directions

DRAW THE DIRECTION OF THE DIPOLES FOR THE
FOLLOWING MOLECULES:
ION-DIPOLE FORCE

Exists between ion and the partial
charge on the end of a polar molecule
Cations attracted to negative end of
dipole
 Anions attracted to positive end of dipole


Magnitude of attraction increases
as charge of ion increases
 as magnitude of dipole increases


Stronger than dipole-dipole forces
DIPOLE-DIPOLE FORCES
Not as strong as ion-dipole forces
 Strength of attraction increases with increasing
polarity

Example: HCl
LONDON DISPERSION FORCES

The constant motion of electrons sometimes
creates temporary dipoles for an instant
Momentary uneven charge creates a dipole
 Can induce a dipole in another molecule

LONDON DISPERSION FORCES

Weakest intermolecular force

Sometimes called dipole induced dipole
All molecules have LDF
 Dependent on motion of electrons so more electrons
means more chances for LDF



LDF strength increases
with increasing molar
mass
Example: O2
VALENCE-BOND THEORY
Covalent bonding involves sharing of electrons
 Electrons exist in orbitals



Orbitals overlap so that electrons can form pairs to
make a bond
As orbitals overlap, they mix and form new
hybrid orbitals

Mixing of the orbitals is called hybridization
HYBRIDIZATION = the blending of orbitals
+
Poodle
+
=
Labrador
+
s orbital
+
=
Labradoodle
=
p orbital
=
sp orbital
EXAMPLE: CH4

Orbital notation for carbon:
Carbon only has 2 electrons available to bond, but it has
to make 4 bonds

Carbon promotes one of its electrons to the 2p orbital
so that each electron can pair up with the 1s electron
from each of the four carbons
1s orbitals of four
hydrogen atoms
Promoted
electrons in carbon
allow 4 bonds
EXAMPLE: CH4

The overlap of carbon’s 1 electron in the s orbital
and 3 electrons in the p orbital creates four
hybrid orbitals

This hybridization is called sp3

The new hybrid orbitals have more energy than an s
orbital but less than a p orbital
HYBRID ORBITALS

To determine hybridization
Draw Lewis structure
 Determine electron domains for target atom
 Hybrid orbitals correlate to number of domains

Orbital Name
Orbitals Combined
Electron domains
sp
1 s / 1 p orbital
2
sp2
1 s / 2 p orbitals
3
sp3
1 s / 3 p orbitals
4
sp3d
1 s / 3 p / 1 d orbital
5
sp3d2
1 s / 3 p / 2 d orbitals
6
SIGMA AND PI BONDS

Sigma () bonds
exist in the region
directly between two
bonded atoms.


Exist on the line
(internuclear axis)
between two atoms
Sigma bonds
s
s
s
p
Single bonds
p
p
SIGMA AND PI BONDS

Pi () bonds exist in the region above and below a line
drawn between two bonded atoms.

Exist perpendicular to the line (internuclear axis) between two
atoms (double and triple bonds)
Pi bonds
Single bond
1 sigma bond
Double Bond
1 sigma, 1 pi bond
Triple Bond
1 sigma, 2 pi bonds