Chemical
Bonding
What is a Bond?


Force that holds atoms together
Results from the simultaneous attraction of
electrons (-) to the nucleus (+)
Breaking/Forming Bonds

When a bond is broken energy is absorbed


When a bond is formed energy is released



Endothermic
Exothermic
The greater the energy released during the
formation of the bond, the greater its
stability
Stable bonds require a great deal of energy
to break
Lewis Dot Diagrams


Use dots to represent the number of
valence electrons
How to write:



Write the symbol.
Put one dot for each valence electron
Electrons go on the 4 sides, no more than 2 per
side
Dot Diagram Examples:

Draw dot-diagrams for the following
1.
2.
3.
Mg
C
Ne
Dot Diagrams - Ions


For ions, use brackets and place the
charge outside the brackets
Examples:
1.
2.
3.
Na+
O2H+
Octet Rule




Atoms will gain or lose electrons in order to
have a full valence shell – like the nobles
gases
“Take the shortest route”
Metals lose electrons to form positive ions
(Cations)
Nonmetals gain electrons to form negative
ions (Anions)
Exceptions





1st principle energy level only holds 2 electrons
Transition elements can lose valence (s) and inner
(d) electrons – this is why they have multiple
oxidation states
Some atoms may be stable with less than an octet
– many compounds with B
Some atoms may be stable with more than an octet
– elements beyond period 2, especially P and S,
the additional electrons are added to the d sublevel
Molecules with an odd number of electrons – they
will be unstable
Types of Bonds


Ionic - Electrons are transferred from a metal to a
nonmetal
Covalent - Electrons are shared between 2
nonmetals



Polar Covalent – electrons are shared unequally
Nonpolar Covalent – electrons are shared equally
Metallic - Electrons are mobile within a metal,
“Sea of Electrons”
Dog Analogy

Ionic Bonds


Polar Covalent Bonds


Dogs of equal strength
Metallic Bonds


Unevenly matched but willing to share
Nonpolar Covalent Bonds


big greedy dog stealing the other dogs bone
Mellow dogs with plenty of bones to go around
See the Dogs
Identifying Bond Type
Ionic – metal and a nonmetal
 Covalent – 2 nonmetals
 Metallic – metals
OR Use electronegativity differences




Ionic: 1.7 or more
Polar Covalent: 0.5-1.6
Nonpolar Covalent: 0.0-0.4
Identifying Bond Types

Indicate the type of bond present in each:
1.
2.
3.
4.
5.
6.
HCl
CCl4
MgCl2
O2
Hg
H2O
Ionic Bonds
Transfer of 1 or more electrons
from a metal to a nonmetal
Electronegativity difference is ≥ 1.7
Example: Sodium Chloride (NaCl)
Na electron
transferred to Cl
Na X Cl
Monatomic Ions
One atom in an ion
Look at the valence electrons to determine
the charges
Examples: K+, O2-
Polyatomic Ions
More than one atom in the ion
Reference Table E
Charge belongs to the entire ion, not an
individual atom
Within the polyatomic ion the atoms are
held together by covalent bonds
When writing it, place ( ) around the entire
ion, with the charge outside
Examples: (NH4)+, (H3O)+, (CO3)2-
Writing Ionic Formulas
You need an equal amount of positive and
negative charges, so that the compound is
neutral
Ionic Formulas are always written as
empirical formulas (reduced)
Examples
1. Na1+ + Cl12. Mg2+ + Cl13. Ca2+ + CO32-
4. Al3+ + O2-
Criss Cross Method
1. Write the symbol for the cation and anion
2. Write each ion’s charge as a superscript
3. Criss-cross the charges to become
subscripts of the other ion
Do not put (+) or (-) charges in the final
formula
4. Reduce to least common multiple
(empirical formula)
Ionic Formulas
Write the formula for the compound
formed from the following ions:
1.
2.
3.
4.
Mg2+ + ClCa2+ + CO32Al3+ + O2Ca + OH
Naming Ionic Compounds
Name the cation first, the anion second
Cation keeps its name, anion changes its
ending to –ide (Chlorine → Chloride)
Do not change the ending of polyatomic
ions
Examples:
1. NaCl
2. CaCO3
3. MgF2
Stock System –
only used for positive ions
Some cations have more than one
positive oxidation states
A roman numeral is used to indicate the
charge of the positive ion
Stock System Examples
1.
2.
3.
4.
Iron (II) Chloride
Iron (III) Oxide
Copper (II) Oxide
a. What charge does copper have in copper
II sulfate?
b. What is the formula for copper II sulfate?
Ionic Salts
Salts are ionic compounds made up of cations
and anions
The ratio of cations to anions is always such
that an ionic compound has no overall charge
Many of the ions are bonded together to form
a crystal
Properties of Ionic Salts
Ionic Bonds are very strong
Very high melting and boiling points
Hard
Brittle
Properties of Salts (cont’d)
Do not conduct electricity as solids
Do conduct electricity when the salt melts
or is dissolved in water (liquid phase or
aqueous)
– In order to conduct electricity a substance
must have free moving charged particles
– In the solid phase the ions are not free to
move
Melting and Boiling Points of
Compounds
Compound
Name
Formula Type of Compound
mp
(oC)
bp
(oC)
Magnesium Flouride
MgF2
Ionic
1261
2512
Sodium Chloride
NaCl
Ionic
801
1686
Calcium Iodide
CaI2
Ionic
784
1373
Iodine MonoChloride
ICl
Covalent
27
370
Carbon tetrachloride
CCl4
Covalent
-23
350
Hydrogen Flouride
HF
Covalent
-83
293
Hydrogen Sulfide
H2S
Covalent
-86
212
Methane
CH4
Covalent
-182
109
Covalent Bonds
Sharing of electrons between 2 nonmetals
 Electronegativity difference is ≤ 1.6

Non-Polar Covalent

Electrons are shared equally

Uniform distribution of electrons
Bond is symmetrical
Electronegativity difference of 0-0.4
All diatomic molecules have non-polar
covalent bonds



Nonpolar Covalent Examples
1.
Flourine (F2)
a.
b.
2.
e-neg difference =
Dot diagram:
Hydrogen (H2)
a.
b.
e-neg difference =
Dot diagram:
Polar Covalent

Unequal Sharing of electrons

Unequal distribution of electrons



Partial positive and partial negative charges
The side with the higher electronegativity will
have a greater share of the electron(s)
resulting in a partial negative charge
Electronegativity difference of 0.5-1.6
Polar Covalent Examples
1.
HCl
a.
b.
2.
e-neg difference:
Dot diagram:
H 2O
a.
b.
e-neg difference:
Dot diagram:
Dipoles
Form when the charge in a bond is
asymmetrical


Present in polar bonds
Partial positive and partial negative charges
Polar Bonds / Dipoles





Isn’t a whole charge just a partial charge
d+ means a partially positive
d- means a partially negative
Example:
d+ dH - Cl
+---→
The Cl pulls harder on the electrons (more eneg)
The electrons spend more time near the Cl
Dipole Examples
Which molecule contains more polar
bonds?
a. CCl4
b. CH4
2. Which has a stronger dipole?
1.
a.
b.
HCl
HBr
Properties of Molecular Substances
(Covalent Compounds)
Soft
 Low melting points and boiling points


Many exist as gases
Poor conductors of heat and electricity (in
all phases)
Examples: H2O, CCl4, NH3, C6H12O6, O2

Molecular Formulas
(Covalent Compounds)
Contain covalent bonds
 Tells you how many atoms are present in a
single molecule
 Named similarly to ionic compounds,
except use prefixes to indicate the number
of atoms per molecule

Prefixes

Mono-
1
Hexa-
6
Di-
2
Hepta-
7
Tri-
3
Octa-
8
Tetra-
4
Nona-
9
Penta-
5
Deca-
10
Mono- is only used for the second element

Example: CO = carbon monoxide
Examples
1.
2.
3.
4.
5.
CCl4
H 2O
NO
N 2O 5
BBr3
Structural Formulas



Specifies how atoms
are bonded together
Dashes represent
bonds
2 atoms can share up
to 3 pairs of electrons
Single Bonds
2 atoms share 1 pair of electrons (2
electrons)
Examples:
1. Ammonia (NH3)
2. Chlorine (Cl2)
3. Hydrochloric Acid (HCl)

Double Covalent Bonds
2 atoms share 2 pairs of electrons (4
electrons)
 2 bonds between 2 atoms
Examples:
1. Carbon Dioxide (CO2)
2. Oxygen (O2)

Triple Covalent Bond
2 atoms share 3 pairs of electrons (6
electrons)
 3 bonds between 2 atoms
Examples:
1. Nitrogen (N2)
2. Ethyne (C2H2)

Bond Length/Strength

Length:

Single > Double > Triple
 The
more electrons in a bond, the greater the
attraction, therefore shorter

As you move down a group bond length
increases
 Due

to increasing molecular size
Strength:

Triple is the strongest, most stable, requires
the most energy to break
Network Solids

Covalently bonded
atoms are linked into
a giant network
(macromolecules)

Examples: Diamond
(C), Graphite (C),
Silicon Carbide (SiC),
and Silicon Dioxide
(SiO2)
Network Solids

Properties:
Hard
 High melting and boiling points
 Do not conduct heat and electricity

Metallic Bonding

Sea of Electrons

Electrons are free to move through the solid.
+
+ + +
+ + + +
+ + + +
Properties of Metallic Solids
Very Strong
 Good conductors of heat and electricity
because electrons are free to move about
 Luster
 High melting point (except Hg)
 Malleable, Ductile

VSEPR Theory

In a small molecule, the electron pairs are
as far away from each other as possible

VSEPR = Valence Shell Electron Pair
Repulsion
Linear







Drawn on a straight line
All molecules of only 2 atoms are linear
Many 3 atom molecules are linear, if there are no
unshared electron pairs on the central atom
If both ends are the same, the molecule is
nonpolar (Symmetrical = Nonpolar)
If the ends are different, the molecule will be polar
(Asymmetrical = Polar)
Bond Angle = 180o
See Molecules
Examples: H2, CO2, HCl
Trigonal Planar
Trigonal Planar
A central atom is bonded to 3 other atoms,
with no extra electrons on the central atom
 Forms a flat “Y” shape (triangle shape)
 If the ends are all the same, NONPOLAR
 If the ends are different, POLAR
 Bond Angle = 120o

See Molecules

Examples: BCl3, BH2F
Pyramidial
A central atom is bonded to 3 other
atoms and the central atom has an
unshared electron pair
 3-D, like a pyramid
 Always POLAR
 Bond Angle = 107o
See Molecules
 Example: NH3

Tetrahedral
Tetrahedral
A central atom bonded
to 4 other atoms
 3-D shape allows the
electron pairs to get as
far away from each
other as possible

H
H
109.5º
C
H
H
Tetrahedral
If all the ends are the same, NONPOLAR
 If the ends are different, POLAR
 Bond Angle = 109.5o
See Molecules
 Examples:
1. CH4
2. CH3Cl

Bent
Bent
A central atom is bonded to 2 other atoms
and the central atom has 2 unshared
electron pairs
 Always POLAR
 Bond angle = 105o
See Molecules
 Example: H2O

Intermolecular Attractions/Forces


Forces between molecules
Determines boiling point, melting point, vapor
pressure, surface tension


The stronger the intermolecular attractions, the
higher the boiling point
All intermolecular attractions are weaker than
actual bonds
Dipole-Dipole Forces




Occurs between 2 polar molecules
The positive end of one molecule is attracted to
the negative end of another molecule
The greater the electronegativity difference is,
the more polar the bond will be and the stronger
the dipole will be
Example: HCl
Dipole Examples
1. Which would have the strongest intermolecular
forces? Explain Why.
a. HCl
b. HBr
2. Which would have the weakest intermolecular
forces? Explain Why.
a. H2S
b. H2O
Hydrogen Bonds


Special, Strong type of Dipole
Attractions
Attraction of a covalently
bonded H atom to a F, O, or
N atom on another covalent
compound
+
d
d
H O
H d+
Hydrogen Bonds



VERY STRONG
Molecules with H bonds will have high boiling
points, melting points, and surface tension
Example: NH3
H-bonds Examples
1. Which sample has Hydrogen Bonds?
a. H2
b. HF
c. F2
d. HCl
2. Which is the strongest?
a. Hydrogen Bonds
b. Covalent Bonds
c. Dipole-Dipole Attractions
Molecule – Ion Attractions



Attraction between a polar compound and an ion (ionic
salt)
Polar substances (such as water) attract ions from ionic
compounds in solution
This allows ionic substances to dissolve in polar
solvents (water)
The anion is attracted to the positive end of the polar solvent
 The cation is attracted to the negative end of the polar
solvent
 The ion dissociates (falls apart)
Example: NaCl(aq)

Molecule-Ion Examples
1. Molecule-Ion attractions are present in which
sample?
a. HCl(l)
b. HCl(aq)
c. KCl(l)
d. KCl(aq)
2. When sodium chloride dissolves in water the
chloride ion is attracted to
a. The positive part of the water, the O atom
b. The negative part of the water, the O atom
c. The positive part of the water, the H atom
d. The negative part of the water, the H atom
Van Deer Waals Forces




Very weak
Exist between non-polar molecules
Caused by momentary dipoles
Increases as molecular mass increases
VDW Examples
1. Rank in order from weakest to strongest:




Hydrogen Bonds
Covalent Bonds
Van deer Waals Forces
Dipole-Dipole Attractions
2. Which would have the strongest intermolecular
forces?
a. H2
b. Cl2
c. F2
d. Br2