Read Sections 4.6, 4.7, 4.8, 4.9 and 4.10 before viewing the slide show.
Unit 15
Molecular Compounds
•Lewis Structures for Molecular Compounds and Ions
•Polyatomic Ions
•Naming Molecular Compounds
Lewis Structures for Molecular Compounds
and Ions (4.7, 4.10)
•Nonmetals – those to the right of the stair-step line
can react with each other by sharing electrons
•The driving force is similar to that for ionic
compounds – the attainment of an octet of electrons
as summarized in the Octet Rule
•In molecular compounds, though, the atoms do not
form individual ions, they share electrons throughout
their structure
•In this unit, we will look at drawing Lewis structures
representing these molecular compounds as well as
other aspects of these shared electron systems
Key Aspects of Lewis Structures (4.10)
•Key elements of a Lewis structure for a molecular
compound include:
•The number of valence electrons drawn equals the total
number of valence electrons in the compound or ion.
•Electrons are shared in such a way that each atom is
effectively surrounded by eight electrons (two for
hydrogen, other exceptions exist).
•All or some of the electrons around an atom may be
shared with other atoms in the structure
An Approach to Drawing Lewis Structures (4.10)
•Several approaches exist for drawing Lewis
structures – this is just one you might consider:
1. Draw skeletal structure for the atoms. Things that might be helpful
are that hydrogen can only take one bond so it has to be on the
outside; carbon typically has four bonds, nitrogen has three bonds,
and oxygen has two bonds; polyatomic molecules and ions
typically have a central atom surrounded by other atoms.
2. Connect each atom to its neighbor(s) with a single bond which
consists of two electrons. The bond may be characterized by two
electrons (:) or a single line (—) attaching the two atoms.
3. Place electrons around the outside of the atoms so each has eight
total electrons including all of those bonded to it as well as the
nonbonded (or lone pair) electrons.
(Continued on Next Slide)
An Approach to Drawing Lewis Structures
Continued (4.10)
4. Count the number of valence electrons the structure should have.
This is the sum of the valence electrons on all atoms. For ions,
the number of valence electrons is increased by the charge and for
negative ions and reduced by the charge for positive ions.
5. Count the number of valence electrons in your structure.
6. If the valence electron counts match from steps 4 and 5 you are
done. If there are too many valence electrons in your structure,
take away a nonbonded pair and slide a nonbonded pair from an
adjacent atom to make a double chemical bond. Continue this
process until the number of valence electrons in the structure
matches the number that should be present.
Examples of Drawing Lewis Structures (4.10)
Draw the Lewis structure for CCl4
Step
1. Draw the skeletal structure
Progress
Cl
Cl
2. Connect atoms by a single bond
Cl
3. Add electrons to each atom to total eight electrons – notice
the there are two electrons in each bond
..
:Cl
¨
4. Count up the valence electrons you should have 4 for C and 7
each for 4 Cl equals 32 valence electrons
5. Your structure has 32 valence electrons – it is done
C Cl
Cl
Cl
│
— C — Cl
│
Cl
..
:Cl:
│
..
— C — Cl:
│
¨
:Cl:
¨
Examples of Drawing Lewis Structures (4.10)
Draw the Lewis structure for NH3
Step
1. Draw the skeletal structure
Progress
H
H N H
2. Connect atoms by a single bond
H
│
H — N — H
3. Add electrons to each atom to total eight electrons – notice
the there are two electrons in each bond
H
│
H — N — H
¨
4. Count up the valence electrons you should have = 5 for N and
1 each for 3 H equals 8 valence electrons
5. Your structure has 8 valence electrons – it is done
Examples of Drawing Lewis Structures (4.10)
Draw the Lewis structure for SO2
Step
1. Draw the skeletal structure
2. Connect atoms by a single bond
3. Add electrons to each atom to total eight electrons – notice
there are two electrons in each bond
4. Count up the valence electrons you should have = 6 for sulfur
and 6 for each of 2 oxygens = 6 + 2 × 6 = 18 electrons
5. Your structure has 20 valence electrons – it needs to lose two
electrons to get down to the 18 from Step 4. Remove any two
nonbonded electrons and, to maintain the octet, slide two from
nonbonded electrons from an adjacent atom into the bond.
6. Now there is the correct number of valence electrons and
each atom is surrounded by 8. Bonds count as two each and
nonbonded electrons count one each.
Progress
O S O
O — S — O
..
..
..
: O — S — O:
¨
¨
¨
..
: O ═ S — O:
¨
¨
¨
Polyatomic Ions (4.9)
•An ion formed from two or more bonded atoms is called a polyatomic ion.
•These ions will often react as units and typically do not change during the
course of a chemical reaction.
•There is not a wonderfully logical way of determining the names of polyatomic
ions, though there are some patterns.
•Table 4.4 (on page 106 or the next slide) gives a short list of common
polyatomic ions. Don’t panic – you don’t have to know all of these.
•For our purposes, the focus will be on five polyatomic ions:
•ammonium
NH4+
•hydroxide
OH•nitrate
NO3•sulfate
SO42•phosphate
PO43•On the proctored midterm, I will not give you the information above – you will
be expected to know these five. If you need others, I will give you this
information for them.
Some Common Ions
Positive
Li+
Lithium
Cu+
Na+
Sodium
Hg22+
Potassium
Rubidium
Cesium
Hydrogen
Ag+
NH4+
K+
Rb+
+
1+ Cs
+
H
Negative
Copper (I) or
Cuprous
Mercury (I) or
Mercurous
Silver
Ammonium
F-
Fluoride
ClO-
Hypochlorite
Cl-
Chloride
ClO2-
Chlorite
Bromide
Iodide
Nitrate
Nitrite
Acetate
Permanganate
ClO3ClO4IO4OHSCNHCO3-
Cyanide
Hydride
O22-
Chlorate
Perchlorate
Periodate
Hydroxide
Thiocyanate
Hydrogen Carbonate or
Bicarbonate
Peroxide
Oxide
S2-
Sulfide
Sulfate
SO32-
Sulfite
Carbonate
C2O42-
Oxalate
Chromate
Cr2O72- Dichromate
BrI1- NO3
NO2C2H3O2MnO4CNH-
Mg2+ Magnesium
Cu2+
Ca2+
Calcium
Hg2+
Sr2+
Strontium
Sn2+
Barium
Iron (II) or
Ferrous
Cadmium
Zn2+
Pb2+
2+ Ba2+
Fe2+
Cd2+
Cr2+
Al3+
Cr3+
3+
Fe3+
Mn2+
Chromium (II) or Co2+
Chromous
Aluminum
As3+
Chromium (III) or Co3+
Chromic
Iron (III) or
Ferric
Copper (II) or
O2Cupric
Mercury (II) or
SO42Mercuric
Tin (II) or
CO32Stannous
Zinc
2- CrO42Lead (II) or
Plumbous
Manganese (II) or
S2O32Manganous
Cobalt (II) or
Cobaltous
Arsenic
N3Cobalt (III) or
PO43Cobaltic
3AsO43BO33-
Thiosulfate
Nitride
Phosphate
P3PO33-
Phosphide
Phosphite
Arsenate
AsO33-
Arsenite
Borate
Writing Formulas for and Naming Compounds
with Polyatomic Ions (4.9)
•Writing formulas using polyatomic ions works just like it did with binary ionic
compounds – think of the polyatomic ion as one big ion.
•In naming the compounds, the polyatomic ion retains its name – no changing
of the ending or anything like that.
•Examples:
Na+
SO42-
make
Na2SO4
called sodium sulfate
NH4+ and
S2-
make
(NH4)2S
called ammonium sulfide
Ba2+ and
PO43-
make
Ba3(PO4)2
called barium phosphate
and
Naming Binary Molecular Compounds (4.7-4.8)
•Writing formulas for binary molecular compounds is a little trickier than for ionic
compounds since the atoms don’t really have a charge (they are sharing
electrons) and many different compounds can be formed from the same two
elements.
# of atoms
Prefix
•The naming is fairly direct. Prefixes are used
1
mono(see the table to the right) to indicate the number
of atoms of each type in the molecule. The mono2
diprefix is only used in reference to the second atom
3
triin the compound.
4
tetra•Examples:
CO
carbon monoxide
5
pentaCO2
carbon dioxide
6
hexaN2O
dinitrogen monoxide (also
7
heptaknown as nitrous oxide or
laughing gas)
8
octaNO2
nitrogen dioxide
9
nonaN2O4
nitrogen tetroxide
10
decaN2O5
dinitrogen pentoxide
Summary of
Inorganic
Nomenclature
How many different elements
(not atoms – elements) are in
the compound?
Exactly two
More than two
What types of elements are
they?
Metal-nonmetal
(Ionic compound)
Metal keeps its name. For
Groups 1 and 2, Al, Zn, and
Ag no modification is
necessary. For other metals,
use a Roman numeral in
parentheses to indicate the
charge on each ion of the
metal.
Nonmetal changes its ending
to –ide.
Nonmetal-nonmetal
Metal keeps its name. For Groups 1
and 2, Al, Zn, and Ag no
modification is necessary. For other
metals, use a Roman numeral in
parentheses to indicate the charge on
each ion of the metal.
Polyatomic ions retain their names.
(Molecular or covalent compound)
First element retains its name.
Second element switches to –ide
ending.
Each name is preceded by a prefix
indicating the number of its atoms
(di-, tri-, tetra-, etc.). Mono- is only
used for the second element.
GSB Modified 2/1/2010