Chapter 7
Chemical Formulas and Chemical
Compounds
7.1 Chemical Names and
Formulas


Chemical formulas show the relative #’s
of atoms in a chemical compound
Examples

C12H22O11
C = 12 H = 22 O = 11

Pb(NO3)4
Pb = 1 N = 4

(NH4)2CrO4 N = 2 H = 8 Cr = 1 O = 4
O = 12
Chemical Names and Formulas

Naming monatomic ions


End in -ide
Look up the charges on the periodic table

F vs F-1



F Fluorine
F-1 Fluoride
S vs S-2


S Sulfur
S-2 Sulfide
Writing Formulas for Ionic
Compounds

Cations – positive ions (metal)


Anions – negative ions


metal
nonmetal
Charge on compound = 0
Writing Formulas for Ionic
Compounds

Binary ionic compounds
 Rules


if the charges on the ions are the same, drop
‘em
if the charges are different, criss-cross


Same charges –
+1 Cl-1  Na
NaCl
+2 O-2+ Mg
MgO
Different charges+1 S-2- Na S
 Na
2
+2 Cl-1- MgCl
 Mg
2
Sodium Chloride
Magnesium Oxide
Sodium Sulfide
Magnesium Chloride
Writing Formulas for Ionic
Compounds

Ternary Ionic compounds


Metal + (Polyatomic ion)
When naming, do not use the ending –ide

Sodium Nitrate


Sodium Carbonate


Na+1 NO3-1 Na(NO3)
Na+1 CO3-2
Na2(CO3)
Aluminum Nitrate

Al+3 NO3-1
Al(NO3)3
Writing Formulas for Ionic
Compounds

Aluminum Phosphate


Al+3 PO4-3
Aluminum Bihypophosphite
Al+3 HPO2-2

Al(PO4)
Al2 (HPO2)3
Aluminum Carbonate

Al+3 CO3-2
Al2 (CO3)3
Writing Names for Ionic
Compounds



Front name – positive (cation – metal)
Back name – negative (anion – nonmetal)
Binary ionic compounds – composed of only 2
types of elements ( M + NM) – end in -ide




NaCl
MgCl2
Al2O3
NaH
Sodium Chloride
Magnesium Chloride
Aluminum Oxide
Sodium Hydride
The

BIG Lie
Stock System – use Roman Numerals for naming
compounds with metals that have multiple charges
(the transitions!)
Cory
Matthews
Loves
Normal
Name
Symbol
Charge
Stock Name
Charge
Stock Name
Copper
Cu
+1
Cu (I)
+2
Cu (II)
Mercury
Hg
+1
Hg (I)
+2
Hg (II)
Lead
Pb
+2
Pb (II)
+4
Pb (IV)
Sn
+2
Sn (II)
+4
Sn (IV)
Fe
+2
Fe (II)
+3
Fe (III)
Topanga Tin
Intensely Iron
BIG Lie
The

More exceptions to the Lie!
 Ag is always = +1 charge


DON’T write Ag I
Zn always = +2 charge

DON’T write Zn II
Practice!

Sn3N2


AgOH


Lead (II) Carbonate
Zn(OH)2


Silver Hydroxide
PbCO3


Tin (II) Nitride
Zinc Hydroxide
Fe2(SO4)3

Iron (III) Sulfate
More Practice!!

Cu(HSO2)2


CuSO2


Copper (II) Bihyposulfite
Copper (II) Hyposulfite
CuHSO2

Copper (I) Bihyposulfite
Practice!

LiClO3


LiClO2


Calcium Carbonate
Ca(HCO2)2


Lithium Chlorite
CaCO3


Lithium Chlorate
Calcium Bicarbonite
Fe(NO3)3

Iron (III) Nitrate
Writing Names for
Molecular Compounds

Molecular Compounds – covalent compounds


2 nonmetals
To name, we use prefixes
#
Prefix
#
Prefix
1
mono
6
hexa
2
di
7
hepta
3
tri
8
octa
4
tetra
9
nona
5
penta
10
deca
don’t use the prefix
Mono
on the first atom
Writing Names for
Molecular Compounds


Prefix-name prefix-name-ide
CO
Carbon Monoxide


CO2
Carbon Dioxide


PCl3


CBr4


Calcium tetrabromide
N2O5


Phosphorous trichloride
Dinitrogen pentoxide
SF6

Sulfur hexafluoride
Writing Formulas for
Molecular Compounds


The prefixes = the subscripts.
Do NOT look at the charges.

Sulfur Dioxide
SO2

Disulfur Trioxide
S2O3

Dinitrogen pentoxide
N2O5
Naming Acids


Acid - when a solution yields H+ ions in solution
2 types

Binary


H and one other type of atom
ternary (sometimes called oxy)

acids that have H with a polyatomic ion
Naming Binary Acids

Rules

Hydro__(begninng of name)__ic acid
 Ex. HCl


HBr


H 2S


Hydrobromic acid
HF


Hydrochloric acid
Hydrofluoric acid
Hydrosulfuric acid
H 3P

Hydrophosphoric acid
Writing Formulas from Names
for Acids

Do the criss-cross

Ex. Hydronitric acid


H3 N
Hydroiodic acid


H+1 N-3
H+1 I-1
HI
Hydrosulfuric acid

H+1 S-2
H2 S
Naming Ternary Acids


H + polyatomic
Rules



Do NOT start with hydroIf the ending of polyatomic is –ate
If the ending of polyatomic is –ite


Ate/ite
Example:

ic/ous
H2SO4
+
-2 – sulfate
 H and SO4

sulfuric acid
-ic + acid
-ous + acid
Naming Ternary Acids

H2SO3

HClO4

HClO3

HClO2

HClO
Formulas for Ternary Acids

Use the criss-cross method

Nitric acid

Phosphorous acid
7.2 Oxidation Numbers

Since electrons are shared, there is no
definite charge - we assign the more
electronegative element the “apparent”
negative charge - this is known as the
oxidation #


oxidation numbers can also be positive.
oxidation # - a number assigned to an atom
to show the distribution of elements
Oxidation Numbers

Rules




Free elements = 0
 Ex. Mg = O
Ions = charges
 Ex. F = -1
S = -2
Oxygen (0) = -2
 except in peroxides (H2O2)
O = -1
H = +1
 except in metal hydrides (MgH2, NaH)
H = -1
Oxidation Numbers

….Rules

More electronegative atom gets a (-) charge

Ox #’s add up to 0 in compounds

Ox #’s = the charge in polyatomic ions
Oxidation # Practice

FeO (Iron II Oxide)
O
= -2
Fe = ?
-2 + x = 0
x= 2

Fe2O3 (Iron III oxide)
O = -2
Fe = ?
3(-2) + 2x = 0
x= 3
Oxidation # Practice

H2SO4 (Hydrogen Sulfate or Sulfuric Acid)
O = -2
H = +1
S =
SO4-2
X + 4(-2) = -2
X=6
Oxidation # Practice

H2SO3 (Hydrogen sulfite or sulfurous acid)



H = +1
S = +4
O = -2
SO3-2
X + 3(-2) = -2
X=4
Oxidation # Practice

H2Cr2O7 (Hydrogen Dichromate or Dichromic Acid




NO3-1 (Nitrate)



H = +1
Cr = +6
O = -2
N = +4
O = -2
MgH2 (Magnesium hydride)


Mg = +2
H = -1
7.3 Using Chemical Formulas

Step 1 – be able to calculate molar mass (aka –
formula mass, molecular weight, atomic weight,
atomic mass, gram formula weight, etc.)
 Add atomic weights from the periodic table


round to the nearest 10th place
Examples

CH4
12.0 + 4(1.0) = 16.0 g/mol

MgSO4· 7H2O
24.3 + 32.0 + 4(16.0) + 14(1.0) + 7(16.0) =
246.3 g/mol
7.3 Using Chemical Formulas

Step 2 – be able to convert between grams, moles,
particles, and liters
Liters
22.4 L
Mole
Grams
Atoms,
molecules,
particles
Using Chemical Formulas

Convert 32.0 g of CH4 to moles, liters, molecules,
total atoms, atoms of H





Moles
Liters
Molecules
Atoms
Atoms H
Conversions

Moles CH4


32.0 g
1
1 mole
16.0g
2.0moles
Liters

32.0g
1
1 mole
16.0g
22.4L
1mole
44.8 Liters
Conversions

Molecules


1 mole
16.0g
6.022 x 1023molecules 1.20 x 1024molecules
1mole
1 mole
16.0g
6.022 x 1023atoms
1mole
5
1 mole
16.0g
6.022 x 1023atoms
1mole
4
Atoms


32.0g
1
32.0g
1
6.00 x 1024molecules
Atoms H

32.0g
1
4.80 x 1024molecules
Percent Composition


Percentage Composition - every compound has a
certain percentage of each type of atom
(we measure it by mass)
Formula

% composition = mass element X 100 =
mass compound
Practice - % Composition

Calculate % composition if a compound contains 24 g
of Carbon and 64 g of Oxygen

% composition = mass element X 100 =
mass compound



Total mass compound = 24 + 64 = 88 g
% Composition C = 24 x 100 = 27 % C
88
% Composition O = 64 x 100 = 73% O
88
Practice - % Composition





What is the % composition of Ba(OH)2?
Ba = 137.3 g
O = 2 (16.0) = 32.0g
H = 2 (1.0) = 2.0g
Ba(OH)2 = 171.3 g
%Ba = 137.3
171.3
%O = 32.0
171.3
%H = 2.0
171.3
= 80.2%
=18.7%
=1.1%
Practice - % Composition

What is the % composition of C6H12O6




C = 6 (12.0) = 72.0g
H = 12 (1.0) = 12.0g
O = 6 (16.0) = 96.0g
C6H12O6 = 180.0g
40.0%
6.7%
53.3%
7.4 Determining a compound’s empirical
and molecular formula

Empirical formula - the lowest whole
number ratio of atoms in a compound
(simplest formula)

4 rules to find empirical formula




Cross out the % → change to grams
Divide each by own molar mass
Divide by the smallest number
If needed, multiply by 2 or 3 ONLY if a whole number
ratio isn’t the result of step 3
Determining a compound’s empirical
formula

Ex1 Calculate the empirical formula if there
is 52.17 % C, 13.04% H, and 34.78 % O



52.17 g C
1
13.04 g H
1
34.78 g O
1
1 mole = 4.3 = 2
2.17
12.0 g
1 mole = 13.04 = 6
2.17
1.0 g
1 mole = 2.17 = 1
2.17
16.0 g
C2H6O
Determining a compound’s empirical
formula

Ex2 Calculate the empirical formula is there
is 26.56 % K, 35.41 % Cr, and 38.03 % O



26.56 g K
1 mole = .68
.68
39.1 g
35.41 g Cr 1 mole = .68
.68
52.o g
38.03 g O
1 mole = 2.38
.68
16.0 g
= 1 x2
= 1 x2
= 3.5 x2
K2Cr2O7
Determining a compound’s
empirical formula

Ex3 Find the empirical formula if a
sample contains 5.6 g N and 12.8 g O
Determining a compound’s molecular
formula

Rules
 Same steps as empirical formula + 3 more



Find the mass of the empirical compound
Divide this mass by the given molecular weight
Multiply the empirical formula by this number
Determining a compound’s
molecular formula

Find the molecular formula of a compound (MW
= 144.0 g) with 66.67 % C, 11.11 % H, and
22.22 % O

Empirical = C4H8O



Mass empirical = 72.0 g
144.0 g/72.0 g = 2
Molecular Formula = C8H16O2
Determining a compound’s
molecular formula

Find the molecular formula of that compound
that contains 19.80% C, 2.50% H, 66.10% O,
11.60% N and MW = 242.0




Empirical Formula = C2H3O5N
Mass Empirical = 121 g
242 g / 121 g = 2
C4H3O10N2