Chapter 10
Chemical Quantities
Molar Mass
 MOLAR
MASS—# of grams of an element in
one mole of that element.
= atomic mass of an element with the units
of g/mol
 Get it from the periodic table
Molar Mass
Elements
Find the molar
mass on the
periodic table
round to
0.1 g/mol
Compounds
ADD the molar
mass of every
atom present in
the compound
MOLAR MASS PRACTICE
If you know the MOLAR MASS …

You can convert between grams and moles!
Volume
of gas at STP
1 Mole
Representative
Particles
Mass
Mole ↔ Mass Examples

How many moles in 51.0 g of Na3PO4?
Find Molar Mass: (23.0 x 3) + 31.0 + (16.0 x 4)
= 164.0 g/mol
 51.0 g Na3PO4 x 1 mole Na3PO4= 0.311 mol Na3PO4
164 g Na3PO4

 What
is the mass of 0.70 mol of NH4Cl?
Find Molar Mass: 14.0 + (1.0 x 4) + 35.5
= 53.5g/mol
 0.70 mol NH4Cl x 53.5 g NH4Cl = 37.45 g NH4Cl
1 mol NH4Cl

1 mole = Avogadro’s # = 6.02x1023 representative particles
 How big is 6.02x1023? 602 000 000 000 000 000 000 000
A
representative particle: atom, molecule,
or formula unit
1
mole of nitrogen gas (N2) contains
6.02 x 1023 molecules.

Because N2 is a molecule (Covalent)
1
mole of calcium fluoride (CaF2) contains
6.02 x 1023 formula units

Because CaF2 is a formula unit (ionic)
 How
many atoms of He are in 1 mole?
Volume
of gas at
STP
1 Mole
Mass
Representative
Particles
Type of Substance:
Elements
ionic
Covalent
Unit: Grams (g)
Unit:
Atoms
Formula Units (f.un)
Molecules (mocs)
By COUNT: Avogadro's Number
How big is 6.02x1023?
602 000 000 000 000 000 000 000

If you had Avogadro's number of
un-popped popcorn kernels, and
spread them across the United
States of America, the country
would be covered in popcorn to a
depth of over 9 miles.

If we were able to count atoms at
the rate of 10 million per second, it
would take about 2 billion years to
count the atoms in one mole.

A mole of coke cans would cover
the surface of the earth to a depth
of over 200 miles.
Measuring Matter
 Some
units of measurement indicate specific
numbers.


A pair means 2
A dozen means 12
 Knowing
how count, mass, & volume relate
allows you to convert between them.
 If 1 dozen apples = 12 apples, and 1 dozen
apples has a mass of 2.0 kg, What is the mass
of 90 apples?


90/12 = 7.5 dozen
= 15 kg
7.5 dozen x 2.0 kg
What is a Mole?
 When
dealing with tiny particles (atoms,
ions, compounds), the sample size is
usually very large.
 Counting is not practical.
 Just as a dozen eggs represents 12
eggs, a mole (mol) of a substance
represents 6.02 x 1023 representative
particles of that substance.
 The number 6.02 x 1023 is known as
Avogadro’s number
Moles

Determining the number of atoms in a
mole of a compound:
 How many moles are in a representative
particle (formula unit or molecule) of the
substance?
 This can be determined from the formula.

Example: each molecule of CO2 contains 3
atoms: 1 carbon atom and 2 oxygen atoms.

A mole of CO2 contains 1 mole of carbon
atoms and 2 moles of oxygen atoms.
The Mass of a Mole
 When
dealing with atoms, it is often
easier to work with mass.
 Gram atomic mass (gam)—the
atomic mass of an element expressed
in grams.


The atomic mass of carbon is 12 amu
The gam of carbon is 12 g.
 The
gam is the mass of 1 mole of
atoms of any element.
The Mass of a Mole of a Compound
 To
determine the mass of a mole of a
compound you need to know the formula and
the gam of each atom in the compound.

Add the masses of each atom to get the mass of
the compound
 Sulfur


trioxide (SO3)
1 mole of S (32.1g) & 3 moles of O (16.0 g each)
32.1g + 16.0g + 16.0g + 16.0g = 80.1 g
 Gram
molecular mass (gmm)—the mass of
1 mole of a molecular compound
 Gram formula mass (gfm)—the mass of 1
mole of an ionic compound
Sec. 2 Mole-Mass-Volume
 We
learned about gam, gmm, & gfm last
time. We can use 1 broad term to tell the
mass of a substance.
 Molar Mass—the mass (in grams) of 1 mole
of a substance.

Why do we have the 3 terms then? Sometimes
the term molar mass in unclear. What is the
molar mass of oxygen? Do you mean oxygen
gas (O2)? Then the molar mass is 32.0g (2 x
16.0g). Or do you mean oxygen atoms (O)?
Then the molar mass is 16.0 g.
The Volume of a Mole of Gas
 The
volumes of 1 mole of different solid and
liquid substances are not the same.
 The volumes of 1 mole of different gases are
the same under the same conditions.
 To keep things under the same conditions,
gases are measured at standard temperature
and pressure (STP)
 Standard temp is 0°C
 Standard pressure is 101.3kPa (or 1 atm)
 At STP, 1 mole of any gas has a volume of
22.4 L
Converting Between
moles, particles, mass, and volume
Volume
of gas at STP
Note: to convert
between particles,
mass, and volume,
you have to go
through moles.
22.4 L
1 mol
6.02 x 1023 particles
1 mol
Representative
Particles
1 mol
22.4 L
Mole
1 mol
.
6.02 x 1023 particles
Molar mass
1 mol
1 mol .
molar mass
Mass
Sec. 3: Percent Composition &
Chemical Formulas
 Percent
Composition—The relative amount
of each element in a compound
The % of all elements in the compound must
equal 100%
 % mass of element = grams of element x 100
grams of compound
 Or


% mass of element = molar mass of element x 100
molar mass of compound
Example
 An
8.20 g piece of magnesium combines
completely with 5.40 g of oxygen to form a
compound. What is the % composition?
 First—add 8.20 g & 5.40 g to get the mass of
the compound. 8.20 + 5.40 = 13.60 g
 % Mg = (mass of Mg/mass of compound) x 100

%

% Mg = 8.20g/13.60g x 100 = 60.3%
O = (mass of O / mass of compound) x 100
% O = 5.40/13.60 x 100 = 39.7%
 Does

this make sense?
60.3 + 39.7 = 100
Empirical Formulas
 Empirical
formula—gives the lowest wholenumber ratio of the elements in a compound
 An empirical formula may or may not be the
same as a molecular formula.

If the formulas are different, the molecular formula
is a simple multiple of the empirical formula.
 Examples:



The empirical formula for H2O2 is HO
For CO2, the empirical & molecular formula are
the same.
C6H6 and C2H2 have the same empirical formula:
CH
Molecular Formulas
 You
can determine the molecular
formula if you know empirical formula
and molar mass.
 Divide the molar mass by the empirical
formula mass.
 Multiply this number by all subscripts in
the empirical formula to get the
molecular formula.
Example:
 Find
molecular formula with a molar mass of
60.0g and empirical formula of CH4N
 1st find the empirical formula mass


1 C, 4 H, 1 N
12 + (4 x 1) + 14 = 30 g
 Then,

divide molar mass by empirical mass
60.0g / 30 g = 2
 Multiply

each element subscript by the this.
C2H8N2
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