THE MOLE (a counting unit)……….
➢A mole represents a set or group, much in
the same way that a dozen represents a set
of twelve.
• 1 dozen eggs = 12 eggs; 1 mol eggs = 6.022 
1023 eggs
• 1 dozen carbon atoms = 12 carbon atoms; 1
mol carbon atoms = 6.022  1023 carbon atoms
• The number “6.022  1023 ” is known as
“Avogadro’s Number”.
THE MOLE (a counting unit)……….
• The number “6.02  1023 ” is known as
“Avogadro’s Number”.
• If you write out Avogadro’s number, it looks
like this:
602,200,000,000,000,000,000,000
THE MOLE (a counting unit)……….
➢Remember:
• 1 mole of atoms = 6.02  1023 atoms
• 1 mole of molecules = 6.02  1023 molecules
• 1 mole of formula units = 6.02  1023 formula
units
• 1 mole of ions = 6.02  1023 ions
➢Atomic Mass vs. Molar Mass:
• Atomic mass = the mass of ONE ATOM
expressed in atomic mass units (amu)
▫ Ex: Oxygen’ s atomic mass = 15.999 amu;
this is the average mass of 1 oxygen atom
• Molar mass = the mass of ONE MOLE of
atoms, molecules, or formula units expressed
in grams
• Ex: Oxygen’s molar mass = 15.999 g; this is
the mass of 6.02X1023 oxygen atoms
THE MOLE (a counting unit)……….
Atomic Mass vs. Molar Mass
Element &
Symbol
Atomic Mass—
Average mass of 1
Atom
Molar Mass—Mass
of 1 mole of atoms
(6.02 X 1023 atoms)
Carbon (C)
12.0 amu
12.0 g C
Helium (He)
4.00 amu
4.00 g He
ACIDS: You need to know these!
Hydrochloric Acid= HCl
Nitric Acid = HNO3
Sulfuric Acid = H2SO4
Carbonic Acid = H2CO3
Phosphoric Acid=H3PO4
Acetic Acid= CH3COOH
I. CALCULATING MOLAR MASS:
1. Phosphoric acid
I. CALCULATING MOLAR MASS:
2. Calcium Nitrate
I. CALCULATING MOLAR MASS:
3. Iron (III) Carbonate
II. FINDING PERCENT COMPOSITION:
• Percent Composition = is the percent by
mass of each element in the compound.
▫ The percent composition is always the same
regardless of the size of the sample.
• Ex: H20 is ALWAYS 11.2 % hydrogen and
88.8% oxygen (by mass)
II. FINDING PERCENT COMPOSITION:
▫ The percent composition of the elements
in a compound adds up to 100%.
II. FINDING PERCENT COMPOSITION:
• Formula Used:
Molar Mass of the element in a compound
Total Molar Mass of the compound
• Or think of it as:
Part
Whole
X 100
X 100
II. FINDING PERCENT COMPOSITION:
• Steps:
1.Find molar mass of each element (parts)
2.Find the total molar mass of the
compound (whole)
3.
Part
X 100
Whole
4. Evaluate your answer (percentages
should add to approximately 100%)
II. FINDING PERCENT COMPOSITION:
• Ex: Find the percent composition of
aluminum sulfate.
II. FINDING PERCENT COMPOSITION:
NOTE: Therefore, a 100g sample of
Al2(SO4)3 would contain
_28
g S, and
56
1_6
gO
g Al,
III. CONVERSIONS BETWEEN GRAMS &
MOLES
• KEY IDEA: It is important to know the
following conversion: The molar mass of
a compound or the mass an element is
equal to 1 mole of that compound or
element.
▫ Ex: 1 mole of carbon is equal to 12
grams of carbon!
▫ Ex: 1 mole of water is equal to 18
grams (molar mass) of water
III. CONVERSIONS BETWEEN GRAMS &
MOLES
• Use a T-chart to convert from what you
know to what you do not know. Use your
conversion factors so units cancel out!
III. CONVERSIONS BETWEEN GRAMS &
MOLES
➢Ex #1: How many moles are represented by
11.5 g of magnesium hydroxide
III. CONVERSIONS BETWEEN GRAMS &
MOLES
➢Ex #2: How many grams of aluminum sulfide are
present in 0.44 moles of aluminum sulfide?
IV. CALCULATION of FORMULAS:
2 types of formulas:
1.Molecular formula (M.F.) = shows the
types and numbers of atoms combined
in a single molecule.
• Ex: C6H12O6 (glucose)
▫Actually has 24 atoms!
IV. CALCULATION of FORMULAS:
2.Empirical formula (E.F.) = a chemical
formula showing the simplest whole
number ratio of atoms in a compound.
• Ex: Reduce glucose C6H12O6 by dividing
all the subscripts by 6 – the empirical
formula would be CH2O
IV. CALCULATION of FORMULAS:
• In the chart below, fill in the empirical
formula given the molecular formula:
Molecular
Formula
Empirical
Formula
H2 O2
HO
H2 O
H2O
C8H16
CH2
C2 H 4
CH2
IV. CALCULATION of FORMULAS:
• Notice 2 things:
1.Some molecular
formulas can’t be
reduced (H2O)
2.Several compounds
can share the same
E.F., but each
compound has its own
unique molecular
formula
Molecular
Formula
Empirical
Formula
H2O2
HO
H2O
H2O
C8H 16
CH2
C2H4
CH2
IV. CALCULATION of FORMULAS:
• WHY DO WE USE THE EMPIRICAL FORMULA
IF IT CONTAINS LESS INFORMATION?
▫ It is easier to experimentally
determine the empirical formula of a
compound than the molecular
formula. So the E.F. is usually found
1st as a step towards finding the
molecular formula.
IV. CALCULATION of FORMULAS:
▫ Steps to finding an EMPIRICAL FORMULA:
1. The amount of each element should be listed in
grams. If analysis of the compound is in percent,
convert to grams: assume a 100 g sample of the
compound.
2. Convert grams of EACH element to moles (T-chart).
3. Divide moles of each element by the smallest
number of moles.
4. Round each answer from “3” to the nearest whole
number and write the empirical formula using
these numbers as subscripts.
EMPIRICAL FORMULAS:
• Ex #1: What is the empirical formula for a compound if a sample
contains 0.9g of calcium and 1.6g of chlorine?
EMPIRICAL FORMULAS:
• Ex # 2: What is the empirical formula for a compound that is 40%
C, 6.71% H, and 53.3% O?
Steps to finding a MOLECULAR FORMULA:
1. If it is not given, determine the empirical
formula of the compound.(Follow all steps
necessary to determine an empirical formula)
2.Calculate molar mass of the empirical formula.
3.Divide the molar mass of the compound by the
molar mass determined from the empirical
formula.
4.Round the answer to “3” to the nearest whole
number and multiply all subscripts in the
empirical formula by this number.
MOLECULAR FORMULAS:
Ex #1: What is the molecular formula of a substance that has an
empirical formula of AgCO2 and a molar mass of 304g/mol?
MOLECULAR FORMULAS:
Ex #2: An unknown compound contains 85.64% carbon and 14.36%
hydrogen. It has a molar mass of 42.08 g/mol.Find its molecular
formula.
HYDRATES
V. HYDRATES:
• Suppose an ionic substance is dissolved in
water. The water can be evaporated to
leave the solid ionic compound.SOME
ionic compounds, however, will form solid
crystals that incorporate water
molecules (where the water molecules
are weakly attracted to the ions in the
crystal structure).
V. HYDRATES:
• Hydrate = an ionic compound containing
water molecules incorporated into its
solid crystal structure
▫ In other words, a hydrate is a
collection of anions (negative ions),
cations (positive ions), and water
molecules
V. HYDRATES:
• Anhydrous salt = the ionic compound
that remains after the water has been
removed (usually by heating) from the
hydrate
▫ Note: salt = ionic compound (usually an
ionic compound that does not contain
OH  ions)
V. HYDRATES:
• The heating of a hydrate to produce an
anhydrous salt can be illustrated as
follows:
hydrate + heat anhydrous salt + water
V. HYDRATES:
• Prefixes are used as coefficients in the
following manner in the chemical
formulas for hydrates.
▫ The chemical formula for magnesium
sulfate heptahydrate is
• 7 moles of H2O are associated with
every 1 mole of MgSO4.
V. HYDRATES:
• The chemical formula for barium hydroxide
octahydrate is
• 8 moles of H2O are associated with every 1
mole of Ba(OH)2.
V. HYDRATES:
• What is the name of the following hydrate,
FePO4 • 4H2O?
V. HYDRATES: MOLAR MASS OF HYDRATES:
• When determining the molar
mass of a hydrate, the molar
mass of the associated water
must be included.
V. HYDRATES: MOLAR MASS OF HYDRATES:
EX #1: Calculate the molar mass of calcium sulfate dihydrate.
V. HYDRATES: % Composition
• PERCENT COMPOSITION of HYDRATES:
• Steps:
1. Find molar mass of water.
2. Find the molar mass of the salt.
3. Add the mass of the water and the salt in
order to get the total mass of the hydrate.
Mass of Water (part)
Mass of Hydrate (whole)
X 100
V. HYDRATES: % Composition
Ex: What percentage of water is found in sodium sulfide nonahydrate
?
V. HYDRATES: Empirical Formula
• Steps:
1. Identify what you know and what you do not know.
2. All amounts should be listed in grams.If analysis is in
percent, convert to grams: assume a 100 g sample.
3. Convert grams of anhydrous salt to moles.
4. Convert grams of water to moles.
5. Divide moles by the smallest number of moles.
6. Round each answer from “5” to the nearest whole
number and these numbers become the
COEFFICIENTS for the salt and water.
V. HYDRATES: Empirical Formula
Ex: Suppose a hydrate contains 9.520 g barium iodide and 0.887g
water. Determine the formula of the hydrate.