Unit V: The Mole Concept
5.1-5.2 – Atomic Mass, Avogrados Hypothesis, and the Mole
(pg. 77-85, Hebden)
Today’s Objectives
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Explain the significance of the mole, including:
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Recognize the significance of relative atomic mass, with
reference to the periodic table
Identify the mole as the unit for counting atoms, molecules, or
ions
Perform calculations involving the mole, including:
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Determine the molar mass of an element or compound
The Mole
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Question: how long would it take to spend a mole of 1
Yuan coins if they were being spent at a rate of 1 billion
coins per second?
What is a mole?
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Atoms are REALLY small!
We can’t work with individual atoms or amu’s (atomic
mass units) in the lab
Why?
Because we can’t see things that small
The Mole
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Instead, we work with samples large enough for us to
see and weigh on a balance using units of grams
This creates a problem….
A pile of atoms big enough for us to see contains billions
of atoms!
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Billions of atoms are hard to keep track of in calculations
So, chemists made up a new unit:
THE
MOLE
The Mole
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Just as a dozen eggs equals 12 eggs, a mole =
602,000,000,000,000,000,000,000 or 6.02x1023
It is equal to that number no matter what kind of
particles you’re talking about
It could represent marbles, pencils, or chicken feet
Usually, the mole deals with atoms and molecules
The mole, whose abbreviation is mol, is the SI base unit
for measuring amount of a pure substance
The Mole
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The mole, as a unit, is only used to count very small items
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It represents a number of items, so, we can know exactly how
many items are in 1 mole
The experimentally determined number a mole is called
is Avogrado’s Number, or 6.02x1023
The term representative particle refers to the species
present in a substance:
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Atoms (most often)
Molecules
Formula units (ions)
Pop Quiz
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1 dozen Mg atoms =
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1 mole Mg atoms =
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6.02x1023 Mg Atoms
1 mole Mg(OH)2 =
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12 Mg atoms
6.02x1023 Mg(OH)2 molecules
1 mole O2 =
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6.02x1023 O2 molecules
How big is a Mole?
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1 Mole of soft drink cans is enough
to cover the surface of the earth to
a depth of over 320 km
If you had Avogrado’s number of
unpopped popcorn kernels, and
spread them across China, the
country would be covered in
popcorn to a depth of over 15 km
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
Mollionaire
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Back to that question: How long would it take to spend a
mole of 1 Yuan coins if they were being spent at a rate of
1 billion per second?
Answer:
¥ 6.02 x 10^23/ ¥1 000 000 000
= 6.02 x 10^14 payments = 6.02 x 10^14 seconds
6.02 x 10^14 seconds/60 = 1.003 x 10^13 minutes
1.003 x 10^13 minutes/60 = 1.672 x 10^11 hours
1.672 x 10^11 hours/24 = 6.968 x 10^9 days
6.968 x 19^9 days/365.25 = 1.908 x 10^7 years
It would take 19 million years!
How gases combine
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Early chemist John Dalton (1766-1844) wondered how much
of a given element would bond (react) with a given amount of
another element
He did not assign an absolute mass for individual atoms of
any given element, but rather assigned an arbitrary (relative)
mass to each element
He assumed that hydrogen was the lightest and assigned
hydrogen a unit mass of 1
Through experimentation, he determined that C was 6 times
heavier than oxygen, so he assigned C a mass of 6
Oxygen was found to have a mass 16 times heavier than
hydrogen, so he assigned O a mass of 16
Using this same process, he was able to determine the
relative masses of all of the elements
John Dalton’s Experiment
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Looked at masses of gases
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11.1g H2 reacted with 88.9g O2
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Interpretation O2 is 8 times heavier (look at PT)
46.7g of N2 reacted with 53.3g O2
42.9g C reacted with 57.1g O2
No real pattern
Joseph Gay-Lussac
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Combined gas
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1L of H2 reacts with 1L Cl2 2L of HCl
1L of N2 reacts with 3L H2 2L of NH3
2L of CO reacts with 1L O2 2L of CO2
Concluded that gases combine in simple volume ratios
But why aren’t the volumes of the reactants and products
equal?
Avogrado’s Hypothesis
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Equal volumes of any gas at standard temperature and pressure
contain the same number of molecules
Example:
 1L of N2 reacts with 3L H2 2L of NH3
 Lets say each volume contains 1 molecule, we could then say:
 1 molecule of N2 reacts with 3 molecules of H2 to form 2
molecules of NH3
 Lets count the atoms to prove this:
 Reactants: 2 nitrogens, 6 hydrogens
 Products: 2 nitrogens, 6 hydrogens
Mass is always conserved in a chemical reaction, volume is
not always conserved in a chemical reaction
Avogrado’s Hypothesis
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Let’s look at the other 2 examples (again assuming each
volume of gas contains 1 molecule):
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1L of H2 reacts with 1L Cl2 2L of HCl
Reactants: 2 hydrogen atoms, 2 Cl atoms
 Products: 2 hydrogen atoms, 2 Cl atoms
 2L of CO reacts with 1L O2 2L of CO2
 Reactants: 2 carbon atoms, 4 oxygen atoms
 Products: 2 carbon atoms, 4 oxygen atoms
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If 2L of H2 reacts with 1L of O2, how many litres of
H2O would be produced?
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4 H, 2 O = 2H2O = 2L H2O
Do exercises 2-5 on p. 78
Who can explain this?
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Avogadro’s Hypothesis
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Equal volumes of any gas at standard temperature
and pressure contain the same number of molecules
This Explains the simple volume ratio for gases
Atomic Mass
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The mass of 1 mole of atoms of an element.
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The mass of one mole of “C” atoms is 12.0g
The mass of one mole of “Ca” atoms is 40.1g
Molar Mass (Molecular Mass)
 The
mass of 1 mole of molecules of
an element or compound
Diatomic Elements
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Some elements are naturally diatomic.
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Remember the “gens”
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H2, O2, N2, F2, Cl2, Br2, I2, At2
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Hydrogen, nitrogen, oxygen, halogens
you must remember these
Special elements
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Sometimes Phosphorus is P
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Sometimes Sulphur is S
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Sometimes P4
Sometimes S8
Assume the rest of the elements are monatomic
Finding the Molar Mass of
Compounds
H2O
= 2(1.0) + 16.0 = 18.0 g/mol
 Ca(NO3)2
= 40.1 + 2(14.0) + 6(16.0)
= 164.1g/mol
 Ammonium phosphate
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(NH4)3PO4
= 3(14.0) +12(1.0) + 31.0 + 4(16.0)
= 149.0 g/mol
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HMWK: p80 #6-7