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Chemistry 11
Chapter 4 - The MOLE
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Relative Atomic Mass
Dalton, concerned with how much one element
could combine with a given amount of element,
hypothesized:
 Molecules are made of “atoms” of
various elements
 If compound B contains twice the
mass of element X as does compound A, then
John Dalton
compound B must contain twice as many atoms
(1766-1844)
of X
 Simple compounds are made up of only one atom of
each of the two elements making up the compound
(this was later proved to be wrong)
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Dalton did not attempt to figure out the mass of an
individual atom of any element. (atoms are SO
small!)
Instead, he simply assigned an arbitrary mass to each
element, assuming that hydrogen was the lightest
element (which had a mass of “1”)
Experiments showed that C was 6x heavier than H, so
C was assigned a mass of “6” (similarly O had a mass
of “16” as it was 16x heavier)
In this way, Dalton was eventually able to
calculate “relative masses”
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Nowadays ...
The standard used for measuring atomic mass is Carbon-12
(instead of hydrogen), since C is a very common element
and thus readily available.
Relative atomic mass is measured in units of Atomic Mass
Units (a.m.u.), which is equal to 1/12 the mass of the C-12
atom
Eg. If C-12 has a mass of 12.0000 amu, then H has a mass
of 1.007825 amu
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Other contributions ...
Joseph Gay-Lussac (1778) experimented with N and
O. He found that for all compounds formed from N
and O, the ratios of volumes of gases used were whole
number ratios:
Eg. 100 mL N2 + 100 mL O2 forms NO (Ratio 1:1)
100 mL N2 + 200 mL O2 forms NO2 (Ratio 1:2)
Joseph Gay-Lussac
Avogadro (1811) then made the assumption that “equal
volumes of different gases, at constant temperature and
pressure, contain the same number of particles” (this is
called Avogadro’s Hypothesis)
In other words, if 1 L of gas A reacts with 1 L of gas B,
then there are exactly as many particles of A present as
B. The molecule formed is AB; if 2 L of gas A reacts
with 1 L of gas B, then there are TWICE as many
atoms of A as B, so A2B will form
Amadeo Avogadro
(1776-1856)
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The MOLE - a unit of measurement
Suppose we want to make FeS molecules in such as way that there wouldn’t
be any extra Fe and S atoms left.
10 Fe atoms + 10 S atoms = 10 FeS molecules
100 Fe atoms + 100 S atoms = 100 FeS molecules
10 000 Fe atoms + 10 000 S atoms = 10 000 FeS molecules
BUT individual atoms are extremely small and so
impossible to count out… (in fact the mass of one Fe atom
is 9.27 x 10-23 g, mass of one S atom = 5.33 x 10-23 g)
So how do we deal with this problem?
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The periodic table shows that:
atomic mass of Fe = 55.8 amu
atomic mass of S = 32.1 amu
SO the masses of Fe and S atoms are in the ratio:
mass Fe = 55.8 amu
mass S
32.1 amu
Provided equal numbers of Fe and S atoms are used, then
the ratio will always be 55.8:32.1
Eg.
mass of 1000 Fe atoms = 1000 x 55.8 = 55.8
mass of 1000 S atoms
1000 x 32.1 32.1
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SO to make FeS we could weigh out and react 55.8 g of Fe with 32.1 g of
S
Hence in chemistry the concept of the MOLE (mol)- a unit of
measurement representing the AMOUNT OF SUBSTANCE
Experimentally, it has been found that ...
1 mol particles= 6.02 x 1023 particles
[particles = atoms, ions, molecules, electrons etc.]
6.02 x 1023 is known as AVOGADRO’S NUMBER (NA)
(named in his honor - he didn’t invent it)
Analogies:
SO
1 pair = 2
1 dozen = 12
1 ream = 500 (usually paper)
1 mol = 6.02 x 1023
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Putting it into use...
If 1 mol = 6.02 x 1023 particles
then 1 mol of C atoms = 6.02 x 1023 C atoms
The MOLAR MASS is the mass of ONE MOLE
of particles in grams
Element
Atomic Mass shown on
Periodic Table
Molar mass of element
C
12.0 amu
12.0 g
Fe
55.8 amu
55.8 g
O
16.0 amu
16.0 g
1 mol C = 12.0 g of C
1 mol Fe = 55.8 g of Fe
1 mol O = 16.0 g O
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