1 mole

1 mole marbles = covers Earth to depth of 50 miles
Counting by weighing
Purpose: Calculate the amount of pennies in the
bag by weighing them. ( You may take the
pennies out of the bag to weigh them, but do
not count them)
Data
Calculations
 Atoms are too small to count so we weigh an
amount (mole) and calculate the number
Avogodro
 The
actual size of the molecule is not important
 Avagadro concluded that equal volumes of gas have
equal number of molecules
 By
the way, nobody really knew what a molecule was at the
time
Some words mean numbers
Pair
 Dozen
 Baker’s dozen
 Gross
 Ream

Mole
1 mole = 6.02 x 1023 representative particles
(atoms)
(molecules)
(formula units)
( the number of carbon-12 atoms in 12.00 g)
Atomic mass units


We measure atoms in atomic mass units
Atomic mass unit = AMU = 1.66x10-27 Kg
 It
is defined as 12 for Carbon – 12
 Carbon 12 has six protons and six neutrons
1 proton=1 amu (1.0078amu)
1 neutron= 1 amu (1.0087 amu)
Chemical Measurements


Atomic Mass
The weighted average of all the mass
numbers for all the isotopes of the atom
(a.m.u.)
Formula Mass
The sum of all the atomic masses for all
atoms in the compound. (a.m.u.)
Calculate the atomic mass or formula
masses






Na
Cl
Br2
NaCl
CO2
Mg(OH)2
22.99 amu
35.45 amu
159.80 amu
58.45 amu
44.01 amu
58.33 amu
Molar Mass
=
=
=
The mass of 1 mole
( 6.02 x 1023) of
representative particles
The atomic mass in g
The formula mass in g
Atom, Molecule and Ions

Representative particles – smallest
particle of that substance
Substance Representative Particle
element
molcular compound
ionic compound
atom
molecule
formula unit
Calculate the Molar mass
Ca
 H2
 KNO3
 (NH4)2S

Remember
1 mole
1 mole
= 6.02 x 1023 r.p.
= molar mass(g)
practice
1.
2.
Determine the number of FU in
.866 moles of AgNO3?
Find the mass of
0.98 moles of CaCl2.
Molar Volume
The volume of 1mole of a gas at
standard temperature – 0oC
standard pressure – 1 atmosphere
(pressure at sea level)
1 mole gas at STP = 22.4 L
Remember
1 mole
1 mole
1 mole
= 6.02 x 1023 r.p.
= molar mass(g)
= 22.4 L gas
Multistep problems
How many fu in 18.9 g NaCl?
 How many L in 4.5 x 1013 molecules
Ne?

Solution= solute + solvent
•
•
•
% mass = g solute
g of solution
PPM
= parts of solute
1,000,000 parts of solution
Molarity = Moles solute
 L of solution
Calculate



1. What is the molarity if 13.5 g NaCl in 451 ml
solution?
2. How many g of KOH are in 3.5 L of a .67 M
solution?
3. How many ml of solution would you need to
have 17.5 g of NaOH in a .35 M solution?
Percent Composition
% element = g element x 100%
g total compound
Formulas


Empirical
simplest whole number ratio
Molecular Formula
actual number of atoms in the
formula
Practice:
A compound has
13.5 g Ca
10.8 g O
.675 g H
What is its empirical formula?
Try this one


What is the empirical formula that is
25.9% N ,
74.1% O
Spark




1. How many moles of HCl are there in 1.00
liters of 2.0 molar HCl?
2. How many milliliters of 13.2 g of HCL of a
3.4 M solution?
3. How many oxygen atoms are in 150.2 mL of
2.00 M H2SO4?
4. How many sulfur atoms are in 158.2 g of
aluminum thiosulfate , Al2 (S2O3)3?
Molecular formula
Set up table:
Empircal
formula
Molar mass
Molecuar
A Chemical Reaction


Balance the following reaction
Hydrogen and oxygen react to produce water
H2 + O 2  H 2 O
2H2 + O2  2H2O
Why Do We Balance an Equation /
Reaction?




Makes it look pretty
I do what I’m told
Has something to do with conservation
Of mass
Gas Has Mass



The balanced equation says I should
mix two hydrogens with one oxygen
And then the reaction should go well
Let’s
See!
So if…….



If the space between molecules is so large
and volume of gas at equal pressure
and equal temperature have the same number of
molecules
•How can we compare things that are not gases?
Mass



We would weigh it. Of course.
But reactions go by numbers of molecules and atoms
Not by mass, which is in grams
I
wonder if there is a conversion factor?
How Can We Measure Atoms, the Darn
Things are Too Small?

We can compare the mass of atoms
 This



was done back in the 1800’s
Hydrogen was found to be the smallest element
Everything was “relative” to Hydrogen
So we can compare elements by using hydrogen as a
standard
Consider This About Atoms

Each atom of hydrogen has 1.0 amu and uranium
has 238 amu.
 Which
has more atoms, one gram of H or one gram of
U?

Calcium has 40. amu per atom and helium has 4
amu per atom.
 Approximately
how many times more atoms are in one
gram of helium than in one gram of calcium?
Wow




So if we can compare masses of atoms to
each other
Then we could weigh them to get the proper
proportion of them in a reaction
Wouldn’t that be convenient?
If only there were a conversion factor!