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Chapter 7
Chemical Formulas
&
Chemical Compounds
DHMO
Chapter 7
Chemical Formulas and Chemical
Compounds
Section 1 Chemical Names and Formulas
Section 2 Oxidation Numbers
Section 3 Using Chemical Formulas
Section 4 Determining Chemical Formulas
Chapter 7
Section 1 Chemical Names and
Formulas
Objectives
• Explain the significance of a chemical formula.
• Determine the formula of an ionic compound
formed between two given ions.
• Name an ionic compound given its formula.
• Using prefixes, name a binary molecular
compound from its formula.
• Write the formula of a binary molecular
compound given its name.
• Chemical Formula – shorthand that uses
symbols to tell the kind and number of
atoms in a molecule or formula unit.
Al2(SO4)3
Subscript 2 refers to 2
aluminum atoms
Subscript 4 refers
to 4 oxygen atoms
in the sulfate ion
Totals for each atom:
Al – 2 atoms
S – 3 atoms
O – 12 atoms
Subscript 3 refers to
everything in the
parentheses
Section 1 Chemical Names and Formulas
Chapter 7
Naming Monatomic Ions
• Monatomic cations are identified simply by
the element’s name.
– examples:
• K+ is called the potassium cation
• Mg2+ is called the magnesium cation
• For monatomic anions, the ending of the
element’s name is dropped, and the ending ide is added to the root name.
– examples:
• F– is called the fluoride anion
• N3– is called the nitride anion
Monatomic Ions – Table 1 pg.221
Chapter 7
Section 1 Chemical Names and
Formulas
Common Monatomic Ions
Chapter 7
Section 1 Chemical Names and Formulas
Common Monatomic Ions
Why no Roman Numeral
For silver or zinc?
Note: these are all cations
What about anions?
Cu2O
CuO
Cuprous oxide
Cupric oxide
Naming Monatomic Ions
Writing and Naming Binary
Ionic Compounds
• Mg+2 and Br-1  MgBr2
Magnesium
Bromide
• Al+3 and O-2  Al2O3 Aluminum Oxide
• Fe+2 and O-2  FeO Iron(II) Oxide
Ferrous Oxide
• Fe+3 and O-2  Fe2O3 Iron(III) Oxide
Ferric Oxide
Polyatomic Ions – Table 2 pg.226
Compounds with Polyatomic Ions
• NH4+1 and Cl-1  NH4Cl
Ammonium Chloride
• Ba+2 and NO3-1  Ba(NO3)2
Barium Nitrate
• Fe+3 and CrO4-2 Fe2(CrO4)3
Iron(III) Chromate
Chapter 7
Section 1 Chemical Names and
Formulas
Naming Binary Ionic Compounds,
Compounds Containing Polyatomic Ions, continued
Sample Problem
Write the formula for tin(IV) sulfate.
Chapter 7
Section 1 Chemical Names and
Formulas
Naming Binary Ionic Compounds, Compounds
Containing Polyatomic Ions, continued
Sample Problem Solution
Write the symbols for the ions side by side.
Write the cation first.
4+
Sn
2
4
SO
Cross over the charges to give
subscripts. Add parentheses around the
polyatomic ion if necessary.
4+
2
2
4 4
Sn (SO )
Chapter 7
Section 1 Chemical Names and
Formulas
Naming Binary Ionic Compounds, Compounds Containing
Polyatomic Ions, continued
Sample Problem Solution, continued
The total positive charge is 2  4+ = 8+.
The total negative charge is 4  2 = 8.
The largest common factor of the subscripts is 2,
so the smallest whole-number ratio of ions in the
compound is 1:2.
The correct formula is therefore
Sn(SO4)2.
Chapter 7
Polyatomic Ions
Polyatomic Ions with Multiple
Oxygens
•
•
•
•
ClOClO2ClO3ClO4-
Hypochlorite
Chlorite
Chlorate
Perchlorate
2 less oxygens
1 less oxygen
Root ion
1more oxygen
Root Polyatomic Ions
CO3-2 NO3-1
SiO4-4 PO4-3
SO4-2 ClO3-1
AsO4-3 SeO4-2 BrO3-1
TeO4-2
IO3-1
Molecular compounds are…
 made
of just nonmetals
 smallest piece is a molecule
 can’t be held together by
opposite charge attraction
 can’t use charges to figure out
how many of each atom (there
are no charges present)
Molecular compounds are easier!
 Ionic
compounds use charges to
determine how many of each.
–You have to figure out charges.
–May need to criss-cross numbers.
compounds: the name
tells you the number of atoms.
 Molecular
– Uses prefixes to tell you the exact
number of each element present!
Prefixes for Naming Covalent Compounds
Naming Binary Molecular Compounds
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
Mono
Di
Tri
Tetra
Penta
Hexa
Hepta
Octa
Nona
Deca
•
•
•
•
•
•
N2O
NO
NO2
N2O3
N2O4
N2O5
dinitrogen monoxide
nitrogen monoxide
nitrogen dioxide
dinitrogen trioxide
dinitrogen tetraoxide
dinitrogen pentaoxide
Acids are…
 Compounds
that give off
hydrogen ions (H1+) when
dissolved in water (the Arrhenius definition)
 Will start the formula with H.
 There will always be some
Hydrogen next to an anion.
 The anion determines the
name.
Acids and Salts
• An acid is a certain type of molecular
compound. Most acids used in the
laboratory are either binary acids or
oxyacids.
– Binary acids are acids that consist of two
elements, usually hydrogen and a halogen.
– Oxyacids are acids that contain hydrogen,
oxygen, and a third element (usually a
nonmetal).
Acids and Salts, continued
• In the laboratory, the term acid usually refers
to a solution in water of an acid compound
rather than the acid itself.
• Many polyatomic ions are produced by the
loss of hydrogen ions from oxyacids.
sulfuric acid
nitric acid
phosphoric acid
H2SO4
HNO3
H3PO4
sulfate
nitrate
phosphate
2
4
SO

3
NO
3
4
PO
Naming Acids
• Binary Acids
– HF
– HCl
– HBr
– HI
Hydrofluoric Acid
Hydrochloric Acid
Hydrobromic Acid
Hydroiodic Acid
Oxyacids
• H2SO4
• HNO3
• H3PO4
– HClO4
– HClO3
– HClO2
– HClO
Sulfuric Acid
Nitric Acid
Phosphoric Acid
Perchloric Acid
Chloric Acid
Chlorous Acid
Hypochlorous Acid
Rules for Naming acids: Name it as a
normal compound first
1) If the anion attached to
hydrogen ends in -ide, put the
prefix hydro- and change -ide to
-ic acid
• HCl - hydrogen ion and chloride
ion = hydrochloric acid
• H2S hydrogen ion and sulfide ion
= hydrosulfuric acid
Naming Acids
•
If the anion has oxygen in it, then it ends in
-ate or -ite
2) change the suffix -ate to -ic acid (use no
prefix)
• Example: HNO3 Hydrogen and nitrate
ions = Nitric acid
3) change the suffix -ite to -ous acid (use no
prefix)
• Example: HNO2 Hydrogen and nitrite
ions = Nitrous acid
2 additional rules
4) If the acid has 1 more oxygen than
the –ic acid, add the prefix pera. HClO3 (Hydrogen Chlorate) is chloric acid
b. HClO4 would be perchloric acid
5) If there is 1 less oxygen than the
-ous acid, add the prefix hypo•
HClO2 (Hydrogen Chlorite) is chlorous acid,
then HClO would be hypochlorous
acid
Practice by naming these:
 HF
 H3 P
 H2SO4
 H2SO3
Writing Acid Formulas – in reverse!
• Hydrogen will be listed first
• The name will tell you the anion
• Be sure the charges cancel out.
• Starts with prefix hydro?- there is
no oxygen, -ide ending for anion
• no prefix hydro?
1) -ate anion comes from –ic ending
2) -ite anion comes from –ous ending
Write formulas for these:
 hydroiodic acid
 acetic acid
 carbonic acid
 hydrobromic acid
Visual Concepts
Salt
Click below to watch the Visual Concept.
Visual Concept
Diatomic Elements –elements that are
always bonded with some other atom even
alone they will bond with themselves
H O N Cl Br I F
O N Cl
Br
I
F
Double Salts – compound that contains two different
metal ions with a nonmetal or negative polyatomic ion
NHCO3 – Sodium Hydrogen Carbonate
KNaSe – Potassium Sodium Selenide
Chapter 7
Section 2 Oxidation Numbers
Objectives
• List the rules for assigning oxidation
numbers.
• Give the oxidation number for each
element in the formula of a chemical
compound.
• Name binary molecular compounds using
oxidation numbers and the Stock system.
Oxidation Numbers
• The charges on the ions in an ionic
compound reflect the electron distribution of
the compound.
• In order to indicate the general distribution of
electrons among the bonded atoms in a
molecular compound or a polyatomic ion,
oxidation numbers are assigned to the
atoms composing the compound or ion.
• Unlike ionic charges, oxidation numbers do
not have an exact physical meaning: rather,
they serve as useful “bookkeeping” devices
to help keep track of electrons.
Text’s Rules for Assigning Oxidation Numbers
• In general when assigning oxidation numbers, shared electrons
are assumed to “belong” to the more electronegative atom in
each bond.
• More-specific rules are provided by the following guidelines.
1. The atoms in a pure element have an oxidation number of
zero.
examples: all atoms in sodium, Na, oxygen, O2,
phosphorus, P4, and sulfur, S8, have oxidation
numbers of zero.
2. The more-electronegative element in a binary compound is
assigned a negative number equal to the charge it would have
as an anion. Likewise for the less-electronegative element.
3. Fluorine has an oxidation number of –1 in all of its compounds
because it is the most electronegative element.
4. Oxygen usually has an oxidation number of –2.
Exceptions:
• In peroxides, such as H2O2,
oxygen’s oxidation number is –1.
• In compounds with fluorine, such as OF2, oxygen’s
oxidation number is +2.
5. Hydrogen has an oxidation number of +1 in all compounds
containing elements that are more electronegative than it; it
has an oxidation number of –1 with metals.
6. The algebraic sum of the oxidation numbers of all atoms in an
neutral compound is equal to zero.
7. The algebraic sum of the oxidation numbers of all atoms in a
polyatomic ion is equal to the charge of the ion.
8. Although rules 1 through 7 apply to covalently bonded atoms,
oxidation numbers can also be applied to atoms in ionic
compounds similarly.
Rules for Oxidation Numbers
1. The oxidation number of an element in its elemental form is zero. Examples of
this are N2 (g), O2 (g), Na (s), Cl2 (g), etc.
2. The oxidation number of a monatomic ion is exactly the same as its charge. So,
Group IA ions will all have an oxidation number of +1, since they all lose one electron.
Group IIA ions will all have an oxidation number of +2. Aluminum ions only exist as
Al+3 and will have an oxidation number of +3.
3. The oxidation state of oxygen, in most compounds is -2, i.e., it tends to pull 2
shared electrons toward itself. The exceptions are H2O2, hydrogen peroxide and O2-2,
peroxide, when it is -1. In O2 it is zero (see rule 1.).
4. The oxidation state of hydrogen is almost always +1. The exceptions are H2
(oxidation number = zero , rule 1.) and when hydrogen is bonded to metals in binary
compounds, like LiH, when it is -1.
5. Fluorine is always -1. The other halogens are -1, except when bonded to
oxygen (rule 3. gives oxygen an oxidation number of -2, making halogens bonded to
oxygen positive).
6. Oxidation number of the atoms in a compound must add up to the total charge on
that molecule or ion.
Rules for Assigning Oxidation
Numbers
Click below to watch the Visual Concept.
Visual Concept
NaNO3
Na = +1
O = -2
N=?
Chapter 7
Section 2 Oxidation Numbers
Assigning Oxidation Numbers, continued
Sample Problem E
Assign oxidation numbers to each atom in
the following compounds or ions:
a. UF6
b. H2SO4
c.
ClO3
Chapter 7
Section 2 Oxidation Numbers
Assigning Oxidation Numbers, continued
Sample Problem E Solution
a. Place known oxidation numbers above
–1
the appropriate elements.
UF6
Multiply known oxidation numbers by
the appropriate number of atoms and
place the totals underneath the
–1
corresponding elements.
UF6
–6
Chapter 7
Section 2 Oxidation Numbers
Assigning Oxidation Numbers, continued
Sample Problem E Solution, continued
The compound UF6 is molecular. The sum of the
oxidation numbers must equal zero; therefore,
the total of positive oxidation numbers is +6.
–1
UF6
+6 – 6
Divide the total calculated oxidation number by
the appropriate number of atoms. There is only
one uranium atom in the molecule, so it must
have an oxidation number of +6. +6 – 1
UF6
+6 – 6
Chapter 7
Section 2 Oxidation Numbers
Assigning Oxidation Numbers, continued
Sample Problem E Solution, continued
b. Hydrogen has an oxidation number of +1.
Oxygen has an oxidation number of 2.
The sum of the oxidation numbers must equal
zero, and there is only one sulfur atom in each
molecule of H2SO4.
Because (+2) + (8) = 6, the oxidation number
of each sulfur atom must be +6.
1 6
2
H2 SO4
2 6
8
Chapter 7
Section 2 Oxidation Numbers
Assigning Oxidation Numbers, continued
Sample Problem E Solution, continued
c. The total of the oxidation numbers should equal
the overall charge of the anion, 1.
The oxidation number of a single oxygen atom in
the ion is 2.
The total oxidation number due to the three
oxygen atoms is 6.
For the chlorate ion to have a 1 charge, chlorine
must be assigned an oxidation number of +5.
+5 2
ClO°V
3
+5 6
Chapter 7
Section 2 Oxidation Numbers
Using Oxidation Numbers for Formulas and Names
• As shown in the table in the next slide, many
nonmetals can have more than one oxidation number.
• These numbers can sometimes be used in the same
manner as ionic charges to determine formulas.
– example: What is the formula of a binary compound formed
between sulfur and oxygen?
From the common +4 and +6 oxidation states of sulfur, you
could predict that sulfur might form SO2 or SO3.
Both are known compounds.
Chapter 7
Section 2 Oxidation Numbers
Common Oxidation States of Nonmetals
Chapter 7
Section 2 Oxidation Numbers
Using Oxidation Numbers for Formulas and Names,
continued
• Using oxidation numbers, the Stock system,
introduced in the previous section for naming ionic
compounds, can be used as an alternative to the
prefix system for naming binary molecular compounds.
Homework
Pages 251-254
Numbers for first half/ first quiz
4,6,7,10,11,14,15,16,18,21,23,24,25
Numbers for second half/ second quiz
28,31,32,36,38,39,40,44,50