Chapter 7
Chemical Formulas and
Chemical Compounds
Section 1:
Chemical Names & Formulas
• There are literally thousands of chemicals
• Isn’t always best to use common names
for chemicals (calcium carbonate is
limestone, sodium chloride is salt, and
hydrogen oxide is water)
• Common names don’t give information
about chemical composition.
Section 1:
Chemical Names & Formulas
• Significance of chemical formulas:
– Gives relative number of atoms of each kind
of element.
– Subscripts: small numbers to the right that
tell the number of atoms
– If no subscript then it is understood to be 1
– H2SO4
– 2 hydrogens, 1 sulfur, 4 oxygen
Section 1:
Chemical Names & Formulas
• When parentheses are used you must
multiply inside and out.
–Al2(SO4)3
(2 Aluminums, 3 Sulfurs, 12 Oxygens)
Section 1:
Names of Binary Compounds
• Binary compounds are those formed
from only 2 elements.
• To write their formulas the positive ion is
written first and then the negative.
• To name them use the complete name of
the positive ion and add the negative ion
name but change the ending to “-ide.”
(Sulfur becomes sulfide, oxygen becomes
oxide, phosphorous becomes phosphide)
Section 1:
Formulas of Binary Compounds
• To write the formula of a
compound you must consider the
charges and multiply by adding
subscripts so that the overall
charge on the compound is zero.
• Ex: zinc is (2+) and sulfur is (2-) so:
– ZnS
– Name: Zinc Sulfide
Section 1:
Formulas of Binary Compounds
• Ex: zinc is (2+) and iodine is (1-)
so:
2+ 2- ( = 0 )
– Zn 2+ I 1Subscripts1
2
– ZnI2
– Name: Zinc Iodide
Section 1:
Formulas of Binary Compounds
• How do you know the charge?
– Use the valence electrons
– Group 1 = 1+, Group 2 = 2+, 3+, 4±
– Group 15 = 3-, Grp. 16 = 2-, Grp. 17 = 1– May use charge chart (page 205) for
transition metals.
Assignment:
• Worksheet: Writing formulas and names
for binary compounds.
Section 1:
Stock System of Nomenclature
• Some transition metals have more than
one possible charge:
• Ex. Copper: Cu+ and Cu2+
Iron:
Fe2+ and Fe3+
Lead:
Pb+3 and Pb+4
Tin:
Sn+2 and Sn+4
Section 1:
Stock System of Nomenclature
• The charges of these elements must be
represented in the name of the
compounds.
• Charges are provided by using Roman
numerals in the names
• Ex: Iron (II) oxide and Iron (III) oxide
• Formulas: FeO
Fe2O3
Section 1:
Stock System of Nomenclature
• How do you know how to write the formula???
• Iron (II) combines with oxygen
Fe2+ O2-
(charges equal zero so FeO)
• Iron (III) combines with oxygen
Fe3+ O2- (add subscripts and multiply
2
3
to equal zero
Section 1:
Stock System of Nomenclature
• How do you know how to write the name if
you only see the formula???
• CuBr2
• The name is Copper Bromide but is it
Copper (I) Bromide or Copper (II)
Bromide???
Section 1:
Stock System of Nomenclature
Then +1
If charges are
+1
-2 ≠ 0
-1
Cu Br2
Section 1:
Stock System of Nomenclature
Then +2
If charges are
+2
-2 = 0
-1
Cu Br2
Assignment
• Worksheet:
Naming and Writing Formulas for
Compounds Using the Stock System
Section 1:
Naming Binary Molecular Compounds
• Molecular compounds are those in which
the elements are close together on the
periodic table.
• Ex:
– Nitrogen and Oxygen
– Carbon and Oxygen
– Sulfur and Oxygen
– Phosphorus and Chlorine
Section 1:
Naming Binary Molecular Compounds
• Ex: Compounds of Nitrogen and Oxygen
– N2O
– NO
– NO2
– N2O3
– N2O5
• Newer method of naming is to use the
stock system with Roman Numerals.
• Old traditional method uses prefixes.
Section 1:
Naming Binary Molecular Compounds
Prefixes indicate the number of atoms in the
compound
•
•
•
•
•
•
•
•
•
•
1 atom: Mono
2 atoms: Di
3 atoms: Tri
4 atoms: Tetra
5 atoms: Penta
6 atoms: Hexa
7 atoms: Hepta
8 atoms: Octa
9 atoms: Nona
10 atoms:Deca
Section 1:
Naming Binary Molecular Compounds
• The less electronegative element is written first
and is given a prefix only if it has more than one
atom in the formula.
• Next element has a prefix indicating the number
of atoms and ends typically with “ide.”
• Examples:
–
–
–
–
–
N2O
NO
NO2
N2O3
N2O5
Dinitrogen Monoxide
Nitrogen Monoxide
Nitrogen Dioxide
Dinitrogen Trioxide
Dinitrogen Pentoxide
Assignment
• Worksheet:
Naming and Writing Molecular
Compounds Using Prefixes
Section 1:
Compounds with Polyatomic Ions
• Many compounds are composed of
polyatomic ions (a group of covalently
bonded atoms that carry a charge).
• Examples of polyatomic ions:
– Sulfate (SO4)2– Nitrate (NO3) –
– Phosphate (PO4)3– Carbonate (CO3)2– Dichromate (Cr2O7)2– Ammonium (NH4)+
Section 1:
Compounds with Polyatomic Ions
• Most polyatomic ions end with “–ate” or “ite” but there are a few exceptions:
–Cyanide (CN)–Hydroxide (OH)-
Note of caution:
Don’t confuse these with binary
compounds since they end in “ide.”
Section 1:
Naming Compounds with Polyatomic Ions
• Simply write the complete name of the
positive element and the name of the
polyatomic ion.
• KNO3 = Potassium Nitrate
• CaSO4 = Calcium Sulfate
• Al(OH)3 = Aluminum Hydroxide
Section 1:
Writing Compounds with Polyatomic Ions
• Writing the formulas for these compounds
are a little trickier.
• Make sure that you treat the polyatomic ion
as a whole unit and do not change its
subscripts!
• (SO4)2- = 1 sulfate ion
• (SO4)2-2 = 2 sulfate ions NOT…
(S2O8)
Section 1:
Writing Compounds with Polyatomic Ions
• Examples:
• Potassium nitrate
Totals: 1+
Charges: +
Symbols: K
Final Formula:
and 1= 0
(NO3)
KNO3
(no parenthesis needed since only 1 ion is required
Section 1:
Writing Compounds with Polyatomic Ions
• Examples:
• Aluminum Sulfate
Totals:
6+
3+
Charges:
Symbols: Al
Add Subscripts:
2
Final Formula:
and
(SO4)
62-
= 0
3
Al2(SO4)3
(parenthesis must be used to show 3 sulfate ions)
Section 1:
Writing Compounds with Polyatomic Ions
• Polyatomic ions may be paired with
transition metals that have multiple
charges.
• Ex: Copper (II) and sulfate = CuSO4
• But Copper (I) and sulfate = Cu2SO4
• When naming them the Roman numeral
must be included.
• Fe3(PO4)2 = Iron (II) Phosphate
Assignment:
• Worksheet:
Naming and Writing Formulas for
Compounds Containing Polyatomic Ions
Section 1:
Naming Acids and Salts
• Memorize the formulas for the common
acids.
• All begin with one or more H atoms.
– Sulfuric Acid H2SO4
– Hydrochloric Acid HCl
– Nitric Acid HNO3
– Phosphoric Acid H3PO4
– Carbonic Acid
H2CO3
Section 1:
Naming Acids and Salts
• Binary acids contain only 2 elements
– Example: Hydrochloric acid HCl
• Oxyacids contain hydrogen, oxygen,
and one other element
– Example: Sulfuric acid H2SO4
Section 1:
Naming Acids and Salts
• When acids have less oxygen atoms
than normal the names change:
– Normal HClO3 is chloric acid
– Loss of 1 oxygen atom HClO2 is
chlorous acid
– Loss of 2 oxygen atoms HClO is
hypochlorous acid
– An extra oxygen atom HClO4 is
perchloric acid
Section 1:
Naming Salts
• Any ionic compound composed of a
cation and the anion from an acid is
referred to as a salt.
• Example:
– NaCl (anion from hydrochloric acid)
– CaSO4 (anion from sulfuric acid)
Section 2
Oxidation Numbers
Section 2: Oxidation Numbers
• Oxidation numbers are numbers assigned
to the atoms in a molecular compound or
ion that indicates the general distribution of
electrons among bonded atoms.
• Oxidation numbers are not actual charges.
• Oxidation numbers are useful in naming
compounds and writing formulas.
+2
-1
+1
+3
-2
Section 2: Oxidation Numbers
• Rules for assigning oxidation numbers:
– Atoms in a pure element have an
oxidation number of zero – O2 Ox. # = 0
– Fluorine always has ox. # of -1
– Oxygen almost always has ox. # of -2
except in peroxides such as H2O2 – then
it is a -1.
Section 2: Oxidation Numbers
• (Rules continued):
– Hydrogen’s ox. # is +1 unless it is with
metals – then it is -1
– The sum of the ox. # in molecules must
be zero, but in polyatomic ions, it is equal
to the ions charge.
Section 2: Oxidation Numbers
• What are the oxidation numbers for each
atom in these compounds?
UF6 :
Fluorine is -1 x 6 = -6
Uranium +6
{+6 + (-6)} = 0
H2SO4 :
Oxygen is -2 (x 4 = -8)
Hydrogen is +1 (x 2 = +2) so
Sulfur has to be +6
{ (+6) + (+2) + (-8) }= 0
Section 2: Oxidation Numbers
• What are the oxidation numbers for the
chlorate polyatomic ion?
ClO3- : Oxygen is -2 x 3 = -6
Chlorine must be +5
{ (+5) + (-6)} = -1 (the ion’s charge)
Section 2: Oxidation Numbers
• Assignment:
– Page 219, question 1, A-K
Section 3:
Using Chemical Formulas
Section 3: Using Chemical Formulas
• With a chemical formula, you can
calculate many characteristic values for
a compound.
• Formula Mass:
– Compounds have masses – just like
elements.
Section 3: Using Chemical Formulas
• Formula Mass:
– The formula mass of any molecule,
formula unit, or ion is the sum of the
average atomic masses of all the atoms
represented in its formula.
– To find the mass of a compound simply
add the masses of the atoms that make
up the compound. Units are amu’s.
Section 3: Using Chemical Formulas
• To find the formula mass of sulfuric acid
(H2SO4):
element # of atoms x mass (to 2 decimals)
H
2
1.01 = 2.02 amu
S
1
32.01=32.01 amu
O
4
16.00=64.00 amu
98.03 amu
Section 3: Using Chemical Formulas
• To find the formula mass of Calcium Nitrate
Ca(NO3)2
element # of atoms x mass =
Ca
1
40.08 =40.08 amu
N
2
14.01 =28.02 amu
O
6
16.00= 96.00 amu
164.10 amu
Section 3: Using Chemical Formulas
• Molar Mass
– The mass of a mole of any substance is
equal to its formula mass – except instead
of amu’s it is in grams.
– Formula mass of sulfuric acid = 98.03 amu
– Molar mass of sulfuric acid = 98.03 grams
Percentage Composition
• It is sometimes useful to know what the
percentage of a compound is an element.
• What percentage of water is oxygen?
H: 1.01 x 2 = 2.02
O: 16.0 x 1 = 16.0
Molar Mass = 18.02 g
16.0 ÷18.02
= 88.79%
Section 3: Using Chemical Formulas
• Molar Mass can be used as a
conversion factor.
1 mole H2SO4
or
98.03 grams
98.03 grams
1 mole H2SO4
Section 3: Using Chemical Formulas
• How many moles are there in 25 g of
H2SO4?
25 g H2SO4 x 1 mole H2SO4 =
98.03 grams
0.255 mol
Section 3: Using Chemical Formulas
• What is the mass of 4.2 moles of
H2SO4?
4.2 mol H2SO4
x
98.03 g H2SO4
1 mol H2SO4
=
411.73 mol
Section 3: Using Chemical Formulas
• How many molecules are in 54 g of
H2SO4?
54 g H2SO4
x
6.02 x 1023 molecules H2SO4
=
98.03 g H2SO4
3.32 x 1023
molecules
Section 4
Determining
Chemical Formulas