Chemical Formulas

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Chapter 7

• Chemical Formulas
• Chemical Compounds

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Important Chemistry in the News Link below
DHMO
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Chemical Formulas and Chemical Compounds
Chapter 7
half
Section 1 Chemical Names and Formulas
Section 2 Oxidation Numbers Section 3
Using Chemical Formulas Section 4 Determining
Chemical Formulas
QUIZ
half
QUIZ
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Objectives
Section 1 Chemical Names and Formulas
Chapter 7

• 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.

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• Chemical Formula shorthand that uses symbols to
  tell the kind and number of atoms in a molecule
  or formula unit.

Al2(SO4)3
Subscript 4 refers to 4 oxygen atoms in the
sulfate ion
Subscript 2 refers to 2 aluminum atoms
Subscript 3 refers to everything in the
parentheses
Totals for each atom Al 2 atoms S 3
atoms O 12 atoms
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Section 1 Chemical Names and Formulas
Chapter 7

• Naming Monatomic Ions
• Monatomic cations are identified simply by the
  elements name.
• examples
• K is called the potassium cation
• Mg2 is called the magnesium cation
• For monatomic anions, the ending of the elements
  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

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Monatomic Ions Table 1 pg.221
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Common Monatomic Ions
Section 1 Chemical Names and Formulas
Chapter 7
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Common Monatomic Ions
Chapter 7
Section 1 Chemical Names and Formulas
Why no Roman Numeral For silver or zinc?
Note these are all cations What about anions?
CuO
Cu2O
Cuprous oxide
Cupric oxide
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(No Transcript)
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Naming Monatomic Ions
x-ide
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Writing and Naming Binary Ionic Compounds

• Mg2 and Br-1 ? MgBr2
  Magnesium Bromide
• Al3 and O-2 ? Al2O3 Aluminum Oxide
• Fe2 and O-2 ? FeO Iron(II) Oxide
  Ferrous Oxide
• Fe3 and O-2 ? Fe2O3 Iron(III) Oxide
  Ferric Oxide

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Polyatomic Ions Table 2 pg.226
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Compounds with Polyatomic Ions

• NH41 and Cl-1 ? NH4Cl Ammonium Chloride
• Ba2 and NO3-1 ? Ba(NO3)2 Barium Nitrate
• Fe3 and CrO4-2? Fe2(CrO4)3 Iron(III)
  Chromate

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Naming Binary Ionic Compounds, Compounds
Containing Polyatomic Ions, continued
Section 1 Chemical Names and Formulas
Chapter 7

• Sample Problem
• Write the formula for tin(IV) sulfate.

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Naming Binary Ionic Compounds, Compounds
Containing Polyatomic Ions, continued
Section 1 Chemical Names and Formulas
Chapter 7

• Sample Problem Solution
• Write the symbols for the ions side by side.
  Write the cation first.

Cross over the charges to give subscripts. Add
parentheses around the polyatomic ion if
necessary.
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Naming Binary Ionic Compounds, Compounds
Containing Polyatomic Ions, continued
Section 1 Chemical Names and Formulas
Chapter 7

• 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 12.
• The correct formula is therefore

Sn(SO4)2.
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Polyatomic Ions
Chapter 7
when 2 cases possible
ite ate
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Polyatomic Ions with Multiple Oxygens

• ClO- Hypochlorite 2 less oxygens
• ClO2- Chlorite 1 less oxygen
• ClO3- Chlorate Root ion
• ClO4- Perchlorate 1more oxygen

hypo - ite - ate - per
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Root Polyatomic Ions
NO3-1
CO3-2
ClO3-1
SO4-2
PO4-3
SiO4-4
BrO3-1
SeO4-2
AsO4-3
IO3-1
TeO4-2
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Molecular compounds are

• made of just nonmetals
• smallest piece is a molecule
• cant be held together by opposite charge
  attraction
• cant use charges to figure out how many of each
  atom (there are no charges present)

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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.
• Molecular compounds the name tells you the
  number of atoms.
• Uses prefixes to tell you the exact number of
  each element present!

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Prefixes for Naming Covalent Compounds
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Naming Binary Molecular Compounds

• Mono
• Di
• Tri
• Tetra
• Penta
• Hexa
• Hepta
• Octa
• Nona
• Deca

• N2O dinitrogen monoxide
• NO nitrogen monoxide
• NO2 nitrogen dioxide
• N2O3 dinitrogen trioxide
• N2O4 dinitrogen tetraoxide
• N2O5 dinitrogen pentaoxide

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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.

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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).

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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 H2SO4 sulfate
nitric acid HNO3 nitrate
phosphoric acid H3PO4 phosphate
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Naming Acids

• Binary Acids
• HF Hydrofluoric Acid
• HCl Hydrochloric Acid
• HBr Hydrobromic Acid
• HI Hydroiodic Acid

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Oxyacids

• H2SO4 Sulfuric Acid
• HNO3 Nitric Acid
• H3PO4 Phosphoric Acid
• HClO4 Perchloric Acid
• HClO3 Chloric Acid
• HClO2 Chlorous Acid
• HClO Hypochlorous Acid

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Rules for Naming acids Name it as a normal
compound first

• 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

NOT
In Book
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Naming Acids

• If the anion has oxygen in it, then it ends in
  -ate or -ite
• change the suffix -ate to -ic acid (use no
  prefix)
• Example HNO3 Hydrogen and nitrate ions Nitric
  acid
• change the suffix -ite to -ous acid (use no
  prefix)
• Example HNO2 Hydrogen and nitrite ions
  Nitrous acid

In Book
NOT
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2 additional rules

• If the acid has 1 more oxygen than the ic acid,
  add the prefix per-
• HClO3 (Hydrogen Chlorate) is chloric acid
• HClO4 would be perchloric acid
• 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

In Book
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Practice by naming these

• HF
• H3P
• H2SO4
• H2SO3

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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?
• -ate anion comes from ic ending
• -ite anion comes from ous ending

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Write formulas for these

• hydroiodic acid
• acetic acid
• carbonic acid
• hydrobromic acid

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Salt
Visual Concepts
Click below to watch the Visual Concept.
Visual Concept
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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
H-7
Double Salts compound that contains two
different metal ions with a nonmetal or negative
polyatomic ion NHCO3 Sodium Hydrogen
Carbonate KNaSe Potassium Sodium Selenide
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Objectives
Section 2 Oxidation Numbers
Chapter 7

• 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.

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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.

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Texts 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.
• 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.

• 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.
• Fluorine has an oxidation number of 1 in all of
  its compounds because it is the most
  electronegative element.

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• Oxygen usually has an oxidation number of 2.
• Exceptions
• In peroxides, such as H2O2, oxygens oxidation
  number is 1.
• In compounds with fluorine, such as OF2, oxygens
  oxidation number is 2.
• 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.

• The algebraic sum of the oxidation numbers of all
  atoms in an neutral compound is equal to zero.
• The algebraic sum of the oxidation numbers of all
  atoms in a polyatomic ion is equal to the charge
  of the ion.
• Although rules 1 through 7 apply to covalently
  bonded atoms, oxidation numbers can also be
  applied to atoms in ionic compounds similarly.

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Bowman's
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 Al3 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.
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Rules for Assigning Oxidation Numbers
Click below to watch the Visual Concept.
Visual Concept
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NaNO3 Na 1 O -2 N ?
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Assigning Oxidation Numbers, continued
Section 2 Oxidation Numbers
Chapter 7

• Sample Problem E
• Assign oxidation numbers to each atom in the
  following compounds or ions
• a. UF6
• b. H2SO4
• c.

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Assigning Oxidation Numbers, continued
Section 2 Oxidation Numbers
Chapter 7

• Sample Problem E Solution
• a. Place known oxidation numbers above the
  appropriate elements.

Multiply known oxidation numbers by the
appropriate number of atoms and place the totals
underneath the corresponding elements.
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Assigning Oxidation Numbers, continued
Section 2 Oxidation Numbers
Chapter 7

• 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.

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.
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Assigning Oxidation Numbers, continued
Section 2 Oxidation Numbers
Chapter 7

• Sample Problem E Solution, continued
• 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.

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Assigning Oxidation Numbers, continued
Section 2 Oxidation Numbers
Chapter 7

• Sample Problem E Solution, continued
• 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
5 ?6
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Using Oxidation Numbers for Formulas and Names
Chapter 7
Section 2 Oxidation Numbers

• 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.

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Common Oxidation States of Nonmetals
Section 2 Oxidation Numbers
Chapter 7
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Using Oxidation Numbers for Formulas and Names,
continued
Section 2 Oxidation Numbers
Chapter 7

• 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.

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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