Chemistry for Earth Scientists
DM Sherman, University of Bristol
2011/2012
Atomic Theory: Elements,
Compounds and Stoichiometry
Chemistry for Earth Scientists
DM Sherman
University of Bristol
Compounds and Chemical Reactions…!
By the 18th Century, it was realized that many substances
were compounds formed by the combination of two or
more elements.
Etienne Francois Geoffroy
1672-1713
Page 1
Chemistry for Earth Scientists
DM Sherman, University of Bristol
2011/2012
The Elements…as of 1808!
The Elements…as of 1937!
Page 2
Chemistry for Earth Scientists
DM Sherman, University of Bristol
2011/2012
18th Century Empirical Observations I!
The Law of conservation of Mass:
“There is no detectable change in mass in an ordinary
chemical reaction.”
+
1.000 g Iron
=
0.432 g Oxygen
1.432 g Hematite
18th Century Empirical Observations II!
The Law of Constant Composition:
“A compound always contains the same elements in the
same proportions by mass.”
=
1.000 g Hematite
+
0.301 g Oxygen
0.698 g Iron
Note: for some compounds/minerals, this is not true!
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Chemistry for Earth Scientists
DM Sherman, University of Bristol
2011/2012
18th Century Empirical Observations III!
The Law of Multiple Proportions:
“The masses of one element that combine with a fixed
mass of the second element are in a ratio of small whole
numbers.”
Mineral
Formula Mass of
Iron (g)
Mass of
Oxygen (g)
Wusite
FeO
1.000
0.286 g
Magnetite
Fe3O4
1.000
0.382
Hematite
Fe2O3
1.000
0.430
4/3
3/2
9/8
Dalton’s Atomic Theory (1808)!
Dalton explained these observations using atomic theory:
•  Elements consist of tiny indivisible
particles called atoms.
•  Atoms of the same element are alike in
mass and size.
•  Atoms combine to form compounds in
simple numerical ratios, such as 1:2, 2:3,
etc.
•  Atoms of two elements may combine in
different ratios to form more than one
compound.
Page 4
Chemistry for Earth Scientists
DM Sherman, University of Bristol
2011/2012
Why is this so important?!
Once we recognize that chemical compounds are
made up of atoms, we can now express reactions
that form compounds in terms of this idea:
Fe + O = FeO (wustite)
2Fe+ 3O = Fe2O3 (hematite)
3Fe + 4O = Fe3O4 (magnetite)
This makes everything much, much simpler! We
can then predict the masses of things produced/
consumed by a chemical reaction if we know the
relative atomic masses.
Law of Combining Volumes!
• In 1808, Gay-Lussac proposed that
gases, under equal conditions of
temperature and pressure, react with
one another in volume ratios of small
whole numbers.
• Amadeo Avogadro (1811) hypothesized
that equal volumes of gasses at the
same P,T contain equal numbers of
molecules (even if they are different
elements!).
Page 5
Chemistry for Earth Scientists
DM Sherman, University of Bristol
2011/2012
The Mole!
•  Dalton had established a system of relative
atomic masses based on hydrogen (relative
mass = 1).
•  Using Avagadro’s hypothesis, Stanislao
Cannizzaro (1861) was able to produce a
system of relative atomic masses by comparing
the relative masses of equal volumes of gas.
•  The original definition of a mole is the number
of oxygen atoms in 16.000 grams of oxygen.
Relative Atomic Mass!
With the invention of mass spectrometry (1912), it was
discovered that atoms of a particular element may have
several different masses! These are the isotopes.
78Kr
80Kr
82Kr
83Kr
84Kr
86Kr
0.356 %
2.27 %
11.6 %
11.5 %
57.0 %
17.3 %
Page 6
Chemistry for Earth Scientists
DM Sherman, University of Bristol
2011/2012
Relative Atomic Mass!
By convention, we take the isotope 12C to be the
standard: an atom of 12C has a mass of 12 atomic
mass units. The masses of other isotopes are given
relative to this. The masses of the elements are the
isotopic masses weighted by their isotopic
abundances.
e.g., Iron = 5.80% 54Fe + 91.72% 56Fe + 2.20% 57Fe
+ .28% 58Fe
This gives a relative atomic mass of 55.85 g/mole.
Masses of Elementary Particles!
As we will learn, atoms consist nuclei
(made up of protons and neutrons)
surrounded by electrons. The isotopes
of a particular element differ in the
number of neutrons.
Isotopic masses still don’t come out to be exact integers
because the masses of the protons and neutrons are not
exactly equal:
• proton mass = 1.67262158 × 10-27 kg
• electron mass = 9.10938188 × 10-31 kg
• neutron mass = 1.67492729(28)×10−27 kg
Page 7
Chemistry for Earth Scientists
DM Sherman, University of Bristol
2011/2012
Determination of Avogadro Constant!
• Faraday (1834) had determined the charge of a mole
of electrons (F = 96485. Coulombs/mole) from
electrochemical reactions.
• Mulliken (1910) had measured the charge of the
electron e =1.602×10−19 C.
• Given the electron charge, we can now work out how
many particles are in a mole:
NA = F/e = 6.022 x 1023/mole
We call this Avogadro’s number.
Molar Mass and Relative Molar Mass!
• Relative molar mass is a dimensionless
quantity:
RMM = Mass of 1 mole/(Mass of 1 mole of 12C)
• Molar mass is the mass (in grams) of 1 mole of
something.
Page 8
Chemistry for Earth Scientists
DM Sherman, University of Bristol
2011/2012
Moles/Mass and Chemical Calculations!
Moles of A
Mass of A
Moles of B
?
Mass of B
The “Unit Factor Method”!
This is how you do chemical calculations. It can also
be applied to many other problems. When you do any
calculation:
• Always write out the units for each quantity.
• Treat the units like algebraic quantities.
• Make sure units cancel to give the desired units for
the answer.
This is, by far, the most important skill I can teach
you!!
Page 9
Chemistry for Earth Scientists
DM Sherman, University of Bristol
2011/2012
Moles/Mass and Chemical Calculations!
Example 1:
Consider the reaction:
CH2O + O2 -> CO2 + H2O
How many grams of CO2 will be produced from
1 gram of CH2O?
Solution:



CH O 1 mole CO
44 g CO


(1g CH O )1mole
30 g CH O  1 mole CH O 1 mole CO
2
2
2
2
2
2


2
= 1.47g CO2
€
Moles/Mass and Chemical Calculations!
Example 2:
How many grams of Cu will be produced by smelting 1
g of chalcopyrite (CuFeS2). The atomic mass of Cu is
63.5 g/mole and the molecular mass of CuFeS2 is
183.6 g/mole.
Solution:




1mole CuFeS
1 mole Cu
63.5 g Cu



(1g CuFeS )183.6
g CuFeS 1mole CuFeS 1 mole Cu 
2
2
2
2
= 0.346 g Cu
€
Page 10
Chemistry for Earth Scientists
DM Sherman, University of Bristol
2011/2012
Important Concepts/Things to Learn!
• The elements
• Conservation of Mass
• The mole
• Elements and Isotopes
• Unit Factor Method
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