Section 6.5 – Molecular Compounds
• Binary molecular compounds are
composed of two nonmetallic atoms.
• Because atoms can combine in
different ratios (for example CO and
CO2 ) we use prefixes to help
distinguish between compounds.
• CO is carbon monoxide
• CO2 is carbon dioxide
• CCl4 is carbon tetrachloride
• Note the –ide ending (similar to how
an anion works, but these aren’t ionic
compounds)
Prefix
Number
mono-
1
di-
2
tri-
3
tetra-
4
penta-
5
hexa-
6
hepta-
7
octa-
8
nona-
9
deca-
10
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Section 6.5 – Molecular Compound Naming
• To convert a name to a formula, write the correct symbols for
the two elements, then add appropriate subscripts.
• If there is just one of the first atom, you don’t need to write the
mono-, it is assumed.
• But for the second atom, if there is one, use mono• Ex: tetraiodine nonoxide is
I4O9
•
sulfur trioxide
SO3
•
phosphorus pentafluoride
PF5
• Ex: N2O
•
PCl3
•
SF6
•
H2O
dinitrogen monoxide
phosphorus trichloride
sulfur hexafluoride
dihydrogen monoxide
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Molecular Bonding - Acids
• Here is a list of some of the most common acids which have
covalent bonds and their names, which don’t always follow the
standard naming convention.
•
•
•
•
•
•
HCl
H2SO4
HNO3
CH3COOH
H3PO4
H2CO3
Hydrochloric acid
Sulfuric acid
Nitric acid
Acetic acid (also written HC2H3O2)
Phosphoric acid
Carbonic acid
• These are the most common ones and good to memorize
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Single Covalent Bonds
• Hydrogen is the simplest model of a covalent bond
• Each Hydrogen has one electron and they share them to
form a single covalent bond.
• The single covalent bond can be represented by the pair
of electrons or as a dash as shown below
•
H:H or H-H
•
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Conventions for naming
• The chemical formulas of ionic compounds describe
_________________________
• (Example: NaCl is a formula unit)
• The chemical formulas of covalent compounds
describe ________________________.
• (Example H2O is a molecule)
• Ionic compounds do not have molecular formulas
because they are not composed of molecules.
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Covalent Molecules
• Combinations of atoms of the nonmetallic elements
in groups 4A, 5A, 6A and 7A of the periodic table
are likely to form ___________________.
• Chemist Gilbert Lewis summarized this tendency in
his formulation of the
_____________________________________
• Sharing of electrons occurs if the atoms
involved acquire the electron configuration of
_______________________.
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Covalent Bonding – Diatomic Gas Fluorine
7
• In a water molecule, two hydrogen atoms form one
single covalent bond each with one oxygen atom.
• Note how the O atom ends up with eight electrons
around it.
• Covalent molecules will form if each atom will end up
with 8 electrons around it (except H).
• Each dot is ______________.
• Each line is ______________.
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Three Single Bonds
Four Single Bonds
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The HONC Rule
 Hydrogen (and Halogens) form _____________________
 Oxygen (and sulfur) form __________________________
 One double bond, or two single bonds
 Nitrogen (and phosphorus) form ____________________
 One triple bond, or three single bonds, or one double bond and
a single bond
 Carbon (and silicon) form _________________________.
 Two double bonds, or four single bonds, or a triple and a single,
or a double and two singles
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Review: The Octet Rule and
Covalent Compounds
 Covalent compounds tend to form so that each
atom, by _____________ electrons, has an
____________ of electrons in its highest
occupied energy level.
 Covalent compounds involve atoms of
_____________________________.

The term “__________________________”
is used exclusively for covalent bonding
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Coordinate Covalent Bonds
• A covalent bond in which ______________________
__________________________________________
__________________________________________.
• This is signified by showing coordinate covalent bonds
as ___________ that point from the atom donating
the pair of electrons to the atom receiving the bond.
• Many polyatomic cations and anions contain both
covalent and coordinate bonds. NH4+ is an example.
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Resonance
 Occurs when more than one valid Lewis structure
can be written for a particular molecule, such as
ozone, below.
 These are resonance structures.
The actual structure is an average or
a blend of the resonance structures.
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Resonance in Benzene, C6H6
Each of these
junctions
represents where
a Carbon is
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Section 16.3: Polar bonds and molecules
• Covalent bonds involve sharing electrons between
____________________________
• Sometimes the sharing is equal and the electron resides
halfway in between the atoms, as in a diatomic gas like
N2,Cl2, etc. This is called a
___________________________________________
• If the bonding electrons are shared unequally, this is
called a
______________________________________________
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Polar Bonds
• The greater the electronegativity value, _________________
_________________________________________________.
A high electronegativity atom is not “stealing” electrons as in
the ionic case, but it is moving them in its direction.
• Consider HCl. Hydrogen has an electronegativity of 2.1 and
Chlorine has an electronegativity of 3.0. These values are
quite different, so the covalent bond in HCl is polar. The
shared electrons are pulled in the direction of Cl, because it is
more electronegative. This can be represented as follows:
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Bond Polarity
Electronegativity Differences and Bond Types
Electronegativity
Difference Range
0.0 – 0.4
0.4 – 1.0
1.0 – 2.0
≥ 2.0
Most probable type of
bond
Nonpolar covalent
Moderately polar
covalent
Very polar covalent
Ionic
Example
H–H
d+ dH – Cl
d+ dH–F
Na+ Cl-
(2.1-2.1=0.0)
(3.0-2.1=0.9)
(4.0-2.1=1.9)
(3.0-0.9=2.1)
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Polar Molecules
• Some molecules have dipoles (aka dipolar
molecules), but their polarities line up in such
a way that they ________________________.
• Carbon dioxide is one such example
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Intermolecular Attractionsvan der Waals forces
• The weakest intermolecular attractions are van der Waals forces.
These consist of two possible types, London dispersion forces and
dipole interactions.
• London dispersion forces, (weakest of all intermolecular
interactions) are caused by the motion of electrons. The strength
of dispersion forces increases as the number of electrons
increases.
• For halogens, which have more e- in their outer shell, the major
attraction between them is dispersion forces.
• These forces are weaker for F and Cl (gases at STP). They are
stronger for Bromine, a liquid at STP, and even stronger for Iodine,
a solid at STP.
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Dipole interaction forces
• The second type of van der Waals force is the dipole
interaction, when polar molecules are attracted to one
another.
• The positive region of one molecule is attracted to the
negative region of another.
HCl molecules
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Hydrogen Bonds
• A hydrogen bond is an attractive force where a hydrogen which is
covalently bonded to a very electronegative atom (meaning the H
has a slight δ+ charge on it) is also weakly bonded to an unshared
electron pair of another atom (pair has δ– charge to it).
• This happens because when H bonds to O, F or N, the very polar
bond leaves the H very electron deficient, with essentially an
exposed nucleus (a proton) with no electrons. The H nucleus is
then attracted to a negatively charged unshared electron pair on
another atom.
• The resulting hydrogen bond is only about 5% of the strength of a
regular covalent bond, but it is still the strongest of the
intermolecular forces.
– This is what causes water to be a liquid at room temperature
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Hydrogen bonds in water
• Hydrogen has valence
e- that are not
shielded from the
nucleus by another layer
of electrons.
• Water has this type of
interaction because the
hydrogens have a
slightly + charge and
the oxygen has a
slightly – charge.
• This relatively strong
interaction is called a
hydrogen bond.
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Intermolecular Forces Summary
• Weakest –
• Middle –
• Strongest –
• But all three are still much weaker than a
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Intermolecular Attractions and Molecular
Properties
• The physical properties of a compound depend on the type of
bonding it has – ____________________________________.
• Here are some comparisons of physical properties
Characteristics of Ionic and Covalent Compounds
Characteristic
Ionic Compound
Covalent Compound
Transfer of one or more
electrons between atoms
Sharing of electron pairs
between atoms
Physical state
Solid
Solid, liquid and gas
Melting point
High (usually > 300 C)
Low (usually < 300 C)
Solubility in water
Usually high
High to low
Electrical conductivity of
aqueous solution
Good conductor
Poor conductor or
nonconducting
Representative unit
Bond formation
Type of elements
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