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
2
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
H· +
Hydrogen
atom
H :
·H
Hydrogen
atom
H:H
Hydrogen
molecule
H
• 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
• Each dash represents a pair of shared electrons.
4
Conventions for naming
• The chemical formulas of ionic compounds describe
formula units
• (Example: NaCl is a formula unit)
• The chemical formulas of covalent compounds
describe molecules.
• (Example H2O is a molecule)
• Ionic compounds do not have molecular formulas
because they are not composed of molecules.
• What does that mean?
• Example: Ionic copper(II) oxide is composed of
equal numbers of Cu2+ and O2- ions in a crystal
lattice.
• The formula unit shows the lowest whole-number
ratio of Cu2+ to O2-, which is 1:1, CuO.
5
Ionic versus Molecular Compounds
6
Covalent Molecules
• Combinations of atoms of the nonmetallic elements
in groups 4A, 5A, 6A and 7A of the periodic table
are likely to form covalent bonds.
• Chemist Gilbert Lewis summarized this tendency in
his formulation of the octet rule for covalent
bonding:
• Sharing of electrons occurs if the atoms
involved acquire the electron configuration of
noble gases.
7
Covalent Bonding – Diatomic Gas
Fluorine
F
+
F
F
F
or
F
F
8
• 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 one electron.
• Each line is two electrons. 9
• In an ammonia molecule,
NH3, three H atoms form
single covalent bonds with
one N atom.
• Note how the N has eight
electrons around it.
• When you consider the N
and the three H’s, you
can see the 2s and 2p
orbitals are now full with
10
eight electrons.
• A methane molecule has four
carbon-hydrogen bonds. In each
bond, C and H share the 1s e- from
the hydrogen and a 2s or 2p efrom the carbon.
• Normally C would start with
1s22s22p2 configuration, but by
promoting one 2s e- to 2p, resulting
in 1s22s12p3 it can create a stable
octet with the four H atoms.
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Covalent
Bonding
Different bonding models
for methane, CH4. Models
are NOT reality. Each has
its own strengths and
limitations.
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Review: The Octet Rule and
Covalent Compounds
 Covalent compounds tend to form so that each
atom, by sharing electrons, has an octet of
electrons in its highest occupied energy level.
 Covalent compounds involve atoms of
nonmetals only.
 The term “molecule” is used exclusively for
covalent bonding
 Ted-Ed video 3:47 What is the shape of a
molecule?
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A single bond
The Diatomic Fluorine Molecule
F
F
1s
2s
2p
1s
2s
2p
Each has
seven valence
electrons
F F
14
Some double bonds:
Carbon Dioxide
A carbon dioxide molecule has two C=O bonds
15
Note how C and the O’s each have 8 electrons now
An exception to the rule:
The Diatomic Oxygen Molecule
O
O
1s
2s
2p
1s
2s
2p
O O
Oxygen has
six valence
electrons.
You would
think O2 would
form a double
bond by the
looks of it, but
experiments
show its
nonstandard
and has two
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unpaired e-
A triple bond:
The Diatomic Nitrogen Molecule
N
N
1s
2s
2p
1s
2s
2p
Each has five
valence
electrons
N N
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I Bring Clay For Our New House
18
Lewis Dot
Structures
 Lewis structures show how valence electrons
are arranged among atoms in a molecule.
 Lewis structures reflect the central idea that
stability of a compound relates to noble gas
electron configuration (atoms will react if
they can arrange themselves to have 8
electrons around them).
 Shared electron pairs are covalent bonds and
can be represented by two dots (:) or by a
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single line ( - )
The HONC Rule
 Hydrogen (and Halogens) form one covalent
bond
 Oxygen (and sulfur) form two covalent bonds
 One double bond, or two single bonds
 Nitrogen (and phosphorus) form three covalent
bonds
 One triple bond, or three single bonds, or one
double bond and a single bond
 Carbon (and silicon) form four covalent bonds.
 Two double bonds, or four single bonds, or a triple
and a single, or a double and two singles
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Completing a Lewis Structure: CH3Cl
 Draw carbon as the central atom (it wants
the most bonds, 4)
 Add up available valence electrons:
 C = 4, H = (3)(1), Cl = 7
 Join peripheral atoms
Total = 14
H
.. ..
.. ..
..
..
to the central atom with
electron pairs.
H
C
Cl
 Complete octets on
atoms other than
H
hydrogen with the
Check: Final structure
remaining electrons
should have 14 e-
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Coordinate Covalent Bonds
• A covalent bond in which one atom contributes both
bonding electrons is called a coordinate covalent bond.
• This is signified by showing coordinate covalent bonds
as arrows 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.
22
Carbon Monoxide – coordinate
covalent bonding
In a coordinate covalent
compound, one atom
contributes both electrons
of a bonding pair. In
carbon monoxide, which
atom contributes two
electrons in one of the
carbon-oxygen bonds?
:C O:
Triple bond – one of them
23
is a coordinate bond
24
exceptions to octet rule
25
How to
represent
ions
(covalent
bonds within
the ion)
26
Electron dot
structure of
the sulfite ion
(SO32-)
each O has 6
each S has 6
plus 2 extra =
26 eSulfur has to
create one
coordinate
covalent bond
to make this
work.
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Bond Dissociation Energies
• Large amounts of heat are given off when
hydrogen atoms combine to make H2, which implies
that the product is more stable than the
reactants.
• If you try to break H2 apart, it will require a
large amount of energy to do it.
• The same thing is true for Carbon. A typical C-C
single covalent bond has a bond dissociation energy
of 347 kJ.
• Since Carbon forms such strong C-C bonds, that
explains why it’s compounds are so stable.
• See table of bond dissociation energies on the
next page. Note which one is weakest.
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Bond Length and Bond Energy
What trend do you notice?
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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 32
Exceptions to the Octet Rule
NO2 (also known as smog in LA) has one unpaired
electron, so it is an exception to the octet rule.
33
Oxygen: an exception to the octet rule
O O
The measured distance between oxygen atoms indicates
that O2 does have some double bond character. This
suggests that oxygen is a hybrid of the two structures
shown on this page.
O O
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Other exceptions
to the octet rule
F – B – F
F
• BF3 is deficient by 2 e-.
• Sometimes Phosphorus or Sulfur
expand the octet to include 10 or 12
electrons.
– Examples are PCl5 and SF6
P
S
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Section 16.3: Polar bonds and molecules
• Covalent bonds involve sharing electrons between
two atoms.
• 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
nonpolar covalent bond.
• This sharing isn’t always equal, because one atom
may pull harder than the other atom, and then
the electrons will not be in the middle.
• If the bonding electrons are shared unequally,
this is called a polar covalent bond (or just a polar
bond).
36
Polar Bonds
• The greater the electronegativity value, the greater
the ability of an atom to attract electrons to itself.
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:
d+
d-
H – Cl
H - Cl
37
Bond Polarity
Electronegativity Differences and Bond Types
Electronegativity Most probable type Example
Difference Range of bond
0.0 – 0.4
0.4 – 1.0
1.0 – 2.0
Nonpolar covalent
Moderately polar
covalent
Very polar covalent
≥ 2.0
Ionic
H–H
(2.1-2.1=0.0)
d+ dH – Cl
(3.0-2.1=0.9)
d+ dH – F (4.0-2.1=1.9)
Na+ Cl- (3.0-0.9=2.1)
38
Sample Problem 16-4
• Find which type of bond (nonpolar covalent, moderately
polar covalent, very polar covalent or ionic) will form
between each of the following?
A) N (3.0) and H (2.1)
Δ = 0.9, moderately polar covalent
B) F (4.0) and F (4.0)
Δ = 0.0, nonpolar covalent
C) Ca (1.0) and O (3.5)
Δ = 2.5, ionic
D) Al (1.5) and Cl (3.0)
Δ = 1.5, very polar covalent
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40
Polar
Molecules
• The presence of a polar bond in a molecule makes the
entire molecule polar. That means one end of the
molecule is slightly negative and the other end
slightly positive.
• A molecule that has two poles is called a dipolar
molecule or a dipole. If this kind of molecule is
placed in an electric field, the molecules orient
themselves with respect to the positive and negative
41
plates creating the field.
Polar Molecules
• Some molecules have dipoles, but
their polarities line up in such a way
that they cancel out.
• Carbon dioxide is one such example
O = C = O
42
43
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.
44
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
45
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
46
temperature
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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48
49
Hydrogen Bonding videos
• Hydrogen bonding
• http://www.youtube.com/watch?v=lkl5cbfqFRM
• Water drops on a penny
• http://www.youtube.com/watch?v=tv4Jrc06yLA&feature=fvwrel
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Intermolecular Forces Summary
•
•
•
•
Weakest – London dispersion forces
Middle – dipole-dipole interactions
Strongest – hydrogen bonds
But all three are still much weaker
than a covalent bond (max 5% of the
strength of an average covalent bond)
51
Intermolecular Attractions and
Molecular Properties
• The physical properties of a compound depend on the
type of bonding it has – ionic or covalent.
• Here are some comparisons of physical properties
Characteristics of Ionic and Covalent Compounds
Characteristic
Ionic Compound
Covalent Compound
Representative unit
Formula unit
Molecule
Bond formation
Transfer of one or more
electrons between atoms
Sharing of electron pairs
between atoms
Type of elements
Metallic & nonmetallic
Nonmetallic
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 Good conductor
aqueous solution
Poor conductor or
nonconducting
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