Chapter 2
Molecular Substances
& Lewis Structures
Definitions:
An element is the simplest form of matter and
contains only one type of atom.
A compound results from the combination of two or
more elements in a specific and constant ratio.
A molecule is a combination of a specific number
and kind of elements held together by the force of a
covalent chemical bond which cannot be broken
without creating a different substance.
Definitions:
A compound results from the combination of two or
more elements in a specific and constant ratio.
A molecule is a combination of a specific number
and kind of elements held together by the force of a
covalent chemical bond which cannot be broken
without creating a different substance.
Methane; the
combination of one
carbon atom and
four atoms of
hydrogen in a
specific
arrangement.
Water; the
combination of one
oxygen atom and
two atoms of
hydrogen in a
specific
arrangement.
Definitions:
A compound results from the combination of two or
more elements in a specific and constant ratio.
A molecule is a combination of a specific number
and kind of elements held together by the force of a
covalent chemical bond which cannot be broken
without creating a different substance.
Salt (sodium chloride);
consist of the elements
sodium and chlorine, in a
specific and constant ratio,
but is considered a
compound, not a molecule.
Distinguishing Compounds and Molecules:
A rule of thumb is that a compound which contains only
nonmetals and/or semimetals will exist as a
molecule. A covalent compound.
NOTE H
I
1
2
3
4
5
6
7
VIII
II
These are
“nonmetals”.
III IV V VI VII
Distinguishing Compounds and Molecules:
A rule of thumb is that a compound which contains only
nonmetals and/or semimetals will exist as a
molecule. A covalent compound.
I
1
2
3
4
5
6
7
VIII
II
These are
“semimetals”.
III IV V VI VII
Distinguishing Compounds and Molecules:
A rule of thumb is that a compound which contains only
nonmetals and/or semimetals will exist as a
molecule. A covalent compound.
NOTE H
I
1
2
3
4
5
6
7
VIII
II
Recall that these
are “metals”.
III IV V VI VII
Distinguishing Compounds and Molecules:
A rule of thumb 2 is that a compound which contains
metals and nonmetals will NOT exist as molecules.
Sodium chloride (NaCl) is formed from a metal and a
nonmetal and is an ionic compound but not a molecule.
I
1
2
3
4
5
6
7
VIII
II
III IV V VI VII
Distinguishing Compounds and Molecules:
Sodium chloride (NaCl) is formed from a metal and a
nonmetal and is an ionic compound. Methane (CH4) and
water (H2O) are formed from two nonmetals and are
covalent compounds that exist as molecules.
I
1
2
VIII
II
III IV V VI VII
3
4
5
6
7
http://www.animatedsoftware.com/elearning/Periodic%20Table/AnimatedPeriodicTable.swf
The above link is for a very cool periodic table. There are many on the web.
Distinguishing Compounds and Molecules:
Sodium chloride (NaCl) is formed from a metal and a
nonmetal and is an ionic compound. Methane (CH4) and
water are formed from two nonmetals and are covalent
compounds that exist as molecules.
We will see that this distinction really reflects the nature
of the chemical bonding between the elements. In
molecules, the atoms share electrons in covalent
bonds, while in ionic compounds, electrons are
transferred from one atom to another (ionic bonds).
Salt (sodium
chloride); an
example of ionic
bonding.
Methane; an
example of
covalent bonding.
Compounds and Molecules:
Many elements can also exist as molecules. The
simplest of these are the diatomic elements.
Hydrogen
Nitrogen
Oxygen
Fluorine
Chlorine
Bromine
Iodine
H2
N2
O2
F2
Cl2
Br2
I2
Note that all of
these are
nonmetals.
Bromine (Br2); an
example of a
diatomic molecule.
In-Class Problem:
Cl is a nonmetal
I
1
2
VIII
II
III IV V VI VII
Mg is a metal
3
4
5
6
7
Would you expect
MgCl2 to be
molecular?
Not molecular.
In-Class Problem:
F is a nonmetal
I
1
2
VIII
H is a nonmetal
II
III IV V VI VII
3
4
5
6
7
Would you expect
HF to be
molecular?
Molecular.
In-Class Problem:
Cl is a nonmetal
I
1
2
II
Si is a semimetal
III IV V VI VII
3
4
5
6
7
VIII
Would you expect
SiCl4 to be
molecular?
Molecular.
In-Class Problem:
Na is a metal
I
1
2
II
VIII
Br is a nonmetalIII IV V VI VII
3
4
5
6
7
Would you expect
NaBr to be
molecular?
Not Molecular.
In-Class Problem:
I
1
2
VIII
II
S is a nonmetal
III IV V VI VII
3
4
5
6
7
Would you expect
S8 to be
molecular?
Molecular.
Nomenclature of Binary Compounds:
A compound consisting of only two elements is called a
binary compound. To name binary molecular
compounds, you simply list the elements with the
element appearing to the left in the periodic table first.
Add the suffix ide to the root of the second element and
use multipliers if necessary.
Multipliers
one
two
three
four
five
six
seven
eight
nine
ten
mono
di
tri
tetra
penta
hexa
hepta
octa
nona
deca
Nomenclature of Binary Compounds:
A compound consisting of only two elements is called a
binary compound. To name binary molecular
compounds, you simply list both elements, with the
element appearing to the left in the periodic table first.
Add the suffix ide to the root of the second element and
use multipliers if necessary.
Carbon appears to the
left and is first in the
name. Mono- is not
used for the first
element.
Chlorine becomes
chloride, and there
are four of them
(tetra).
Give an acceptable
name for CCl4.
Carbon tetrachloride
Nomenclature of Binary Compounds:
A compound consisting of only two elements is called a
binary compound. To name binary molecular
compounds, you simply list both elements, with the
element appearing to the left in the periodic table first.
Add the suffix ide to the root of the second element and
use multipliers if necessary.
When both elements are
in the same group, the
lower one comes first.
Sulfur is therefore first
in the name.
Oxygen becomes
oxide, and there are
two of them (di).
Give an acceptable
name for SO2.
Sulfur dioxide
Nomenclature of Binary Compounds:
A compound consisting of only two elements is called a
binary compound. To name binary molecular
compounds, you simply list both elements, with the
element appearing to the left in the periodic table first.
Add the suffix ide to the root of the second element and
use multipliers if necessary.
Nitrogen appears to the
left, is first in the name
and there are two (di).
Oxygen becomes
oxide, and there is
only one of them
(mono).
Give an acceptable
name for N2O.
Dinitrogen monoxide
Nomenclature of Binary Compounds:
A compound consisting of only two elements is called a
binary compound. To name binary molecular
compounds, you simply list both elements, with the
element appearing to the left in the periodic table first.
Add the suffix ide to the root of the second element and
use multipliers if necessary.
Nitrogen appears to the
left, is first in the name and
there are two (di).
Oxygen becomes
oxide, and there
are five of them
(pent). “a” is
dropped in front of
oxide.
Pentaoxide is
changed to
pentoxide.
Give an acceptable
name for N2O5.
Dinitrogen pentoxide
In-Class Review:
Indicate whether the following compounds are
molecular.
a.
b.
c.
d.
e.
1
2
3
4
5
6
7
I
NiCl2
H2S
SiH4
SO3
CaH2
II
Not molecular (metal-nonmetal)
Molecular (nonmetal-nonmetal)
Molecular (semimetal-nonmetal)
Molecular (nonmetal-nonmetal)
Not molecular (metal-nonmetal)
VIII
III IV V VI VII
In-Class Review:
Indicate whether the following compounds are
molecular.
a.
b.
c.
d.
e.
1
2
3
4
5
6
7
I
Cs2O2
I2O7
OF2
CH4
NaCl
II
Not molecular (metal-nonmetal)
Molecular (nonmetal-nonmetal)
Molecular (nonmetal-nonmetal)
Molecular (nonmetal-nonmetal)
Not molecular (metal-nonmetal)
VIII
III IV V VI VII
In-Class Review:
Name the following compounds.
a.
b.
c.
d.
e.
1
2
3
4
5
6
7
I
CO2
H2S
SiH4
SO3
NCl3
II
Carbon dioxide
Dihydrogen monosulfide
Silicon tetrahydride
Sulfur trioxide
Nitrogen trichloride
VIII
III IV V VI VII
In-Class Review:
Name the following compounds.
a.
b.
c.
d.
e.
1
2
3
4
5
6
7
I
P4S6
I2O7
OF2
CH4
S4N2
II
Tetraphosphorous hexasulfide
Diiodine heptoxide
Oxygen difluoride
Carbon tetrahydride (methane)
tetrasulfur dinitride
VIII
III IV V VI VII
In-Class Review:
Show formulas for the following compounds.
a.
b.
c.
d.
e.
1
2
3
4
5
6
7
I
Nitrogen triiodide
Diphosphorous pentoxide
Sulfur trichloride
Nitrogen monoxide
Hydrogen monochloride
II
NI3
P2O5
SCl3
NO
HCl
VIII
III IV V VI VII
In-Class Review:
1. Which of the following is/are not a
binary molecular compound?
a.
b.
c.
d.
e.
f.
HCN
N2O5
NH3
HF
BrCl
KCl
In-Class Review:
2. The correct name for the compound P4O6 is:
a.
b.
c.
d.
e.
phosphorous oxide
phosphate ion
phosphorous hexoxide
hexaphosphorous tetroxide
tetraphosphorous hexoxide
In-Class Review:
3. Which of the following is a molecular compound
that contains 7 atoms per molecule?
a.
b.
c.
d.
e.
silicon tetrachloride
sulfur dioxide
dinitrogen pentoxide
carbon dioxide
arsenic trifluoride
In-Class Review:
4. The correct name for the pure, gaseous H2S is:
a.
b.
c.
d.
hydrogen sulfide
dihydrogen sulfide
hydrogen disulfide
dihydrogen monosulfide
In-Class Review:
5. The chemical formula for arsenic pentafluoride is:
a.
b.
c.
d.
e.
ArF5
AsF5
Ar5F
(AsF)5
As5F
Structure of the Atom:
In the previous chapter, we described atomic
structure in terms of a central nucleus surrounded by
a “clouds” of electrons in distinct “shells”. We
differentiated between “core electrons” and the
electrons in the outermost shell, which we termed
valence electrons.
Simple diagram of the
nuclear and electronic
structure of an atom of
carbon.
Four valence
electrons.
Structure of the Atom:
In the previous chapter, we described atomic
structure in terms of a central nucleus surrounded by
a “clouds” of electrons in distinct “shells”. We
differentiated between “core electrons” and the
electrons in the outermost shell, which we termed
valence electrons.
Valence electrons are
present in the outermost
shell and easily move
away from the atom to
form chemical bonds
or ions (charged atoms
or groups).
Four valence
electrons.
Lewis Dot Structures:
In order to help us understand chemical bonding
between atoms, it is often useful to show these
valence electrons using the scheme suggested many
years ago by G.N. Lewis. In this method, valence
electrons are shown as “dots” placed around the
atom. When there are more than four valence
electrons, the electrons are shown as pairs
Using carbon as an
example, you first write
the chemical symbol
and then show all the
valence electrons as
individual dots.
C
Lewis Dot Structures:
When there are more than four valence electrons, the
electrons are shown as pairs. Thus for fluorine, with
seven valence electrons:
F
Lewis Dot Structures:
How do you know how many valence electrons an
element has? The number of valence electrons is
simply given by the group of the element in the
periodic table for the representative elements.
Mg
H
B
I
1
2
3
4
5
6
7
Si
VIII
II
Group
III IV V VI VII
Lewis Dot Structures:
How do you know how many valence electrons an
element has? The number of valence electrons is
simply given by the group of the element in the
periodic table for the representative elements.
P
Se
Cl
I
1
2
3
4
5
6
7
Ne
VIII
II
Group
III IV V VI VII
Lewis Dot Structures:
When there are an unequal number of electrons
and protons in an atom, the atom will be charged
and is called an ion. A positively charged atom or
group is called a cation, a negatively charged atom
or group is called an anion.
When drawing the Lewis structure for an ion, show all
valence electrons and then add or subtract electrons
to give the final charge.
Lewis Dot Structures:
When there are an unequal number of electrons
and protons in an atom, the atom will be charged
and is called an ion. A positively charged atom or
group is called a cation, a negatively charged atom
or group is called an anion.
For chloride anion (Cl-),
there are seven
valence electrons in
chlorine, and one more
to give the negative
charge.
Cl
In-Class Problem:
Draw the Lewis structure for sulfide anion (S2-).
S
2-
I
1
2
3
4
5
6
7
VIII
II
Group
III IV V VI VII
In-Class Problem:
Draw the Lewis structure for nitride anion (N3-).
3-
N
I
1
2
3
4
5
6
7
VIII
II
Group
III IV V VI VII
In-Class Problem:
Draw the Lewis structure for sodium cation (Na+).
Na
I
1
2
3
4
5
6
7
VIII
II
Group
III IV V VI VII
Lewis Dot Structures:
In order to show bonding in most molecular
compounds, atoms share valence electrons so that
each atom has eight valence electrons. This type of
bonding is termed covalent (the octet rule). Eight
valence electrons are present in the noble gases,
making them very stable and unreactive.
In molecular chlorine,
two atoms with seven
valence electrons each
bond to form molecular
chlorine, Cl2.
Cl Cl
Note that each chlorine
atom has eight valence
electrons, including the
shared pair.
Lewis Dot Structures:
In order to show bonding in most molecular
compounds, atoms share valence electrons so that
each atom has eight valence electrons. This type of
bonding is termed covalent (the octet rule). Eight
valence electrons are present in the noble gases,
making them very stable and unreactive.
In molecular chlorine,
two atoms with seven
valence electrons each
bond to form molecular
chlorine, Cl2.
Cl Cl
Note that each chlorine
atom has eight valence
electrons, including the
shared pair.
Lewis Dot Structures:
In order to show bonding in most molecular
compounds, atoms share valence electrons so that
each atom has eight valence electrons. This type of
bonding is termed covalent (the octet rule). Eight
valence electrons are present in the noble gases,
making them very stable and unreactive.
The shared electrons
are often shown as a
dash connecting the
two atoms. This
means an
electron pair.
Cl Cl
Cl Cl
In-Class Problem:
Draw the Lewis structure for molecular nitrogen (N2).
N
N
I
1
2
3
4
5
6
7
Draw the two atoms,
then bond them to give
each atom eight valence
electrons.
VIII
II
Group
III IV V VI VII
In-Class Problem:
Draw the Lewis structure for molecular nitrogen (N2).
Draw the two atoms,
then bond them to give
each atom eight valence
electrons.
I
1
2
3
4
5
6
7
VIII
II
Group
III IV V VI VII
In-Class Problem:
Draw the Lewis structure for molecular nitrogen (N2).
N N
Note that each nitrogen
atom has eight valence
electrons, including the
shared pairs.
Draw the two atoms,
then bond them to give
each atom eight valence
electrons.
In-Class Problem:
Draw the Lewis structure for molecular nitrogen (N2).
N N
Note that each nitrogen
atom has eight valence
electrons, including the
shared pairs.
Draw the two atoms,
then bond them to give
each atom eight valence
electrons.
In-Class Problem:
Draw the Lewis structure for molecular nitrogen (N2).
N N
Note that each nitrogen
atom has eight valence
electrons, including the
shared pairs.
or
N N
a triple bond
In-Class Problem:
Draw the Lewis structure for ammonia (NH3).
H
N
I
1
2
3
4
5
6
7
II
H
H
Draw the CENTRAL
atom, then bond the
other atoms around it to
give each atom eight
valence electrons.
VIII
Group
III IV V VI VII
In-Class Problem:
Draw the Lewis structure for ammonia (NH3).
H
NH
H
Note that the nitrogen
atom has eight valence
electrons, including the
shared pairs and each
hydrogen has two valence
electrons.
Draw the CENTRAL
atom, then bond the
other atoms around it to
give each atom eight
valence electrons.
In-Class Problem:
Draw the Lewis structure for ammonia (NH3).
H
NH
H
Note that the nitrogen
atom has eight valence
electrons, including the
shared pairs and each
hydrogen has two valence
electrons.
H
or
N H
H
a single bond
Two valence electrons (as
in the noble gas He) is a
stable state for hydrogen.
In-Class Problem:
Draw the Lewis structure for carbon monoxide (CO).
C
O
I
1
2
3
4
5
6
7
Draw the two atoms,
then bond them to give
each atom eight valence
electrons.
VIII
II
Group
III IV V VI VII
In-Class Problem:
Draw the Lewis structure for carbon monoxide (CO).
Draw the two atoms,
then bond them to give
each atom eight valence
electrons.
I
1
2
3
4
5
6
7
VIII
II
Group
III IV V VI VII
In-Class Problem:
Draw the Lewis structure for carbon monoxide (CO).
C O
Note that each atom
has eight valence electrons,
including the shared pairs.
Draw the two atoms,
then bond them to give
each atom eight valence
electrons.
In-Class Problem:
Draw the Lewis structure for carbon monoxide (CO).
C O
Note that each atom has
eight valence electrons,
including the shared pairs.
Draw the two atoms,
then bond them to give
each atom eight valence
electrons.
In-Class Problem:
Draw the Lewis structure for carbon monoxide (CO).
C O
Note that each atom has
eight valence electrons,
including the shared pairs.
or
C O
a triple bond
In-Class Problem:
Draw the Lewis structure for sulfur tetrafluoride (SF4).
F
F S F
F
Draw the CENTRAL
atom, then bond the other
atoms around it so that
we can make four
bonds
VIII
I
1
2
3
4
5
6
7
II
Group
III IV V VI VII
In-Class Problem:
Draw the Lewis structure for sulfur tetrafluoride (SF4).
F
F S F
F
We seem to have a problem.
In order to form four bonds,
we must move one electron
pair on sulfur into “storage”
and then split the remaining
pair into single electrons.
Draw the CENTRAL
atom, then bond the
other atoms around it
so that we can make
four bonds
In-Class Problem:
Draw the Lewis structure for sulfur tetrafluoride (SF4).
F
F S F
F
We seem to have a problem.
In order to form four bonds,
we must move one electron
pair on sulfur into “storage”
and then split the remaining
pair into single electrons.
Draw the CENTRAL
atom, then bond the
other atoms around it
so that we can make
four bonds.
“storage”
In-Class Problem:
Draw the Lewis structure for sulfur tetrafluoride (SF4).
F
F S F
F
Draw the CENTRAL
atom, then bond the
other atoms around it
so that we can make
four bonds
Now we can form four
bonds, giving the fluorines
eight electrons each.
“storage”
In-Class Problem:
Draw the Lewis structure for sulfur tetrafluoride (SF4).
F
F S F
F
Now we can form four
bonds, giving the fluorines
eight electrons each.
Finally, we must put our pair of
electrons back on the sulfur.
This gives the sulfur ten
valence electrons. This is
called valence expansion.
Nonmetal elements in period
3 or above may undergo
valence expansion if
needed.
Polyatomic Ions:
A group of atoms covalently bonded together in a fixed
ratio and having a charge is called a polyatomic ion.
In chapter 3, you will learn the names of a number of
these polyatomic ions. Some examples are:
NO3SO42NH4+
OH-
nitrate
sulfate
ammonium
hydroxide
Polyatomic ions act as nonmetals in compounds
because they are molecular or covalently bonded.
In-Class Problem:
Draw the Lewis structure for chlorate anion (ClO3-).
Draw the central atom, add
the other atoms and their
associated electrons.
O Cl O
O
I
1
2
3
4
5
6
7
VIII
II
Group
III IV V VI VII
In-Class Problem:
Draw the Lewis structure for chlorate anion (ClO3-).
Draw the central atom, add
the other atoms and their
associated electrons.
Next add an extra electron
to the chlorine to give the
negative charge.
O Cl O
O
I
1
2
3
4
5
6
7
VIII
II
Group
III IV V VI VII
In-Class Problem:
Draw the Lewis structure for chlorate anion (ClO3-).
O Cl O
O
Group the atoms and their
associated electrons so that each
oxygen has eight electrons.
This also leaves eight electrons
around the chlorine.
I
1
2
3
4
5
6
7
VIII
II
Group
III IV V VI VII
In-Class Problem:
Draw the Lewis structure for chlorate anion (ClO3-).
O Cl O
Or, using line-bond
conventions.
O
I
1
2
3
4
5
6
7
VIII
II
Group
III IV V VI VII
In-Class Problem:
Draw the Lewis structure for molecular nitrogen
monoxide (NO).
N
O
I
1
2
3
4
5
6
7
Draw the two atoms,
then bond them to give
each atom eight valence
electrons.
VIII
II
Group
III IV V VI VII
In-Class Problem:
Draw the Lewis structure for molecular nitrogen
monoxide (NO).
Draw the two atoms,
then bond them to give
each atom eight valence
electrons.
I
1
2
3
4
5
6
7
VIII
II
Group
III IV V VI VII
In-Class Problem:
Draw the Lewis structure for molecular nitrogen
monoxide (NO).
a “free radical”
N O
Note that the oxygen
atom has eight valence
electrons, including the
shared pair, but that
nitrogen only has seven.
or
N O
In-Class Problem:
Draw the Lewis structure for nitrate anion (NO3-).
Draw the central atom, add
the other atoms and their
associated electrons.
O N O
O
I
1
2
3
4
5
6
7
VIII
II
Group
III IV V VI VII
In-Class Problem:
Draw the Lewis structure for nitrate anion (NO3-).
Draw the central atom, add
the other atoms and their
associated electrons.
Next add an extra electron
to the nitrogen to give the
negative charge.
O N O
O
I
1
2
3
4
5
6
7
VIII
II
Group
III IV V VI VII
In-Class Problem:
Draw the Lewis structure for nitrate anion (NO3-).
ONO
O
Group the atoms and
their associated
electrons so that each
oxygen has eight
electrons.
VIII
I
1
2
3
4
5
6
7
II
Group
III IV V VI VII
In-Class Problem:
Draw the Lewis structure for nitrate anion (NO3-).
ONO
O
This is not good, this is
not right, nitrogen only
has six electrons!
I
1
2
3
4
5
6
7
VIII
II
Group
III IV V VI VII
In-Class Problem:
Draw the Lewis structure for nitrate anion (NO3-).
ONO
O
We can fix this, however, by
simply moving a pair of
electrons from one of the
oxygens to form a double
bond to the nitrogen.
I
1
2
3
4
5
6
7
VIII
II
Group
III IV V VI VII
In-Class Problem:
Draw the Lewis structure for nitrate anion (NO3-).
ON O
O
Now all atoms have eight
electrons. Next, let’s
redraw this using dashes
for the covalent bonds.
I
1
2
3
4
5
6
7
VIII
II
Group
III IV V VI VII
In-Class Problem:
Draw the Lewis structure for nitrate anion (NO3-).
O N O
This is a correct
structure for the nitrate
anion...but is it the only
correct structure?
O
I
1
2
3
4
5
6
7
VIII
II
Group
III IV V VI VII
In-Class Problem:
Draw the Lewis structure for nitrate anion (NO3-).
O N O
What if we had chosen
to form the double bond
with the oxygen on the
left?
O
I
1
2
3
4
5
6
7
VIII
II
Group
III IV V VI VII
In-Class Problem:
Draw the Lewis structure for nitrate anion (NO3-).
O N O
How about the oxygen
on the bottom?
O
I
1
2
3
4
5
6
7
VIII
II
Group
III IV V VI VII
In-Class Problem:
Draw the Lewis structure for nitrate anion (NO3-).
O N O
So which one is
correct?
O
All of them!
I
1
2
3
4
5
6
7
VIII
II
Group
III IV V VI VII
In-Class Problem:
O N O
O N O
O
O
These structures are
called resonance forms
and because they differ
only in the arbitrary
placement of electrons,
they are all equivalent.
O N O
O
In-Class Problem:
O N O
O N O
O
O
To indicate that they are
resonance forms, a special
double headed arrow is used to
connect them. It is important to
note that these forms do not
represent distinct structures,
but that the nitrate anion is a
hybrid of all of these forms.
O N O
O
In-Class Problem:
O N O
O N O
O
O
The hybrid looks something like this.
O N O
O N O
O
O
In-Class Problem:
O N O
O
In the hybrid, each
1
oxygen is linked by “1 /
3
bonds” and each
oxygen has
1
/ negative charge.
3
The electron and charge distribution in
the hybrid is actually best described as a
“cloud”, with the unit (the anion) having
one net negative charge.
Another simple example of a resonancestabilized system is the formate anion CHO2-
O
H
O
H
C
O
C
O
The resonance hybrid.
O
H
C
equal bond lengths
O
You should note that resonance forms represent
electronic limits; the resonance hybrid is the ion.
O
H
O
H
C
O
C
O
The resonance hybrid.
O
H
C
equal bond lengths
O
Rules for Drawing Resonance Forms...
• move only electrons, not atoms,
• all resonance structures must be legal
Lewis structures, and finally,
• always remember that resonance forms
are electronic limits; and just like a
mathematical limit, you never get there!
The actual molecule is the hybrid, it
never looks like any of the limits, but is
always something in between.
Using Bonding Patterns in Lewis Structures:
After examining many of the structures we have drawn,
we can reach a few simple conclusions:
¾ Hydrogen and the halogens form one bond and
share two electrons in covalent bonds.
¾ Oxygen usually forms two bonds, shares four
electrons and has two unshared pairs of
electrons.
¾ Nitrogen usually forms three bonds, shares six
electrons and has one unshared pair of electrons.
¾ Carbon usually forms four bonds, shares eight
electrons and has no unshared pairs.
Using Bonding Patterns in Lewis Structures:
After examining many of the structures we have drawn,
we can reach a few simple conclusions:
¾ Hydrogen and the halogens form one bond and
share two electrons in covalent bonds.
H Cl
Using Bonding Patterns in Lewis Structures:
After examining many of the structures we have drawn,
we can reach a few simple conclusions:
¾ Oxygen usually forms two bonds, shares four
electrons and has two unshared pairs of
electrons.
O H
H
Using Bonding Patterns in Lewis Structures:
After examining many of the structures we have drawn,
we can reach a few simple conclusions:
¾ Nitrogen usually forms three bonds, shares six
electrons and has one unshared pair of electrons.
H
N H
H
Using Bonding Patterns in Lewis Structures:
After examining many of the structures we have drawn,
we can reach a few simple conclusions:
¾ Carbon usually forms four bonds, shares eight
electrons and has no unshared pairs.
H
H C H
H
In-Class Problem:
Draw the Lewis structure for hydrazine (N2H4).
H N N H
H H
¾ Draw the nitrogens and their unshared pairs of
electrons.
¾ Each nitrogen will be bonded to two hydrogens.
¾ Make three bonds to each nitrogen by simply
connecting them together.
In-Class Problem:
Draw the Lewis structure for methanol (CH3OH).
H
H C O H
H
¾ Draw the carbon in the middle; it will form four bonds.
¾ The oxygen must form two bonds; attach it to the
carbon.
¾ Put hydrogens all the way around to make the structure.
In-Class Problem:
Draw the Lewis structure for formaldehyde (CH2O).
H C O
H
¾ Draw the carbon in the middle; it will form four bonds.
¾ Put the two hydrogens on the carbon and add the
oxygen.
¾ The oxygen must form two bonds, but because there are
no more atoms, both of these must be to the carbon.
In-Class Problem:
Draw the Lewis structure for formaldehyde (CH2O).
H C O
H
¾ Draw the carbon in the middle; it will form four bonds.
¾ Put the two hydrogens on the carbon.
¾ The oxygen must form two bonds, but because there are
no more atoms, both of these must be to the carbon.