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The History of the Atom
During all of history, there have been many different ideas about what
an atom is (most of them incorrect). In the next unit we’ll discuss how
atomic theory developed historically and figure out what the modern
atom actually looks like.
 Early Greek Models (Leucippus, Democritus, Plato) – 5th
century BC.
o The early Greek philosopher Leucippus was the first guy
who believed that matter was made of tiny, indivisible
particles called atoms. His student Democritus usually
gets credit for this discovery, though.
o The atomic model:
 Matter is made of empty space through which
atoms move.
 Atoms are tiny, indestructible objects.
 Different atoms have different sizes and shapes
(Plato’s addition: These shapes are geometric
solids)
 Changes in matter are due to changes in the
combinations of atoms.
o This theory, along with a non-atomic theory (favored by
Aristotle) reigned for over two millennia.
 Some random observations (18th century)
o For a very long time, nobody did any scientific
experimentation, so the idea of the atom was left in the
realm of pure thought. However, Enlightenment
philosophy taught that reason was the only true basis for
knowing, so experimentation restarted.
o Law of Conservation of Mass (Lavoisier): The weights
of the products in a chemical reaction will be the same as
the weights of the reactants.
 This was a revolutionary idea, because if you burn
wood the ashes weigh less than the wood originally
did.
 Lavoisier was able to figure this out because he did
a lot of work with gases and realized that they
actually have mass.
o Law of definite composition (Proust): A chemical
compound will always have the same composition,
regardless of how it was made.
 This was revolutionary because impurities made it
so this didn’t always seem true.
 John Dalton’s atomic theory (1808):
o All matter is made of small, indestructible particles
called atoms.
 False: Atoms can be broken. A better statement
that’s still true: Matter cannot be broken by
chemical means.
o Atoms of the same element have identical properties.
 False: The presence of isotopes disproves this.
o Atoms of different elements have different properties.
 True
o Atoms obey the law of conservation of mass – the
weight of what we make is equal to the weight of what we
started with.
 True
o Atoms obey the law of multiple proportions (this is
known as Dalton’s law): If two elements can make
more than one chemical compound, the ratios of the
masses of one of the elements that combine with a given
amount of the second element will be a small, whole
number.
 An example: Hydrogen and oxygen make two
chemical compounds. In one compound, 2 grams
of hydrogen combine with 16 grams of oxygen. In
another compound, 2 grams of hydrogen combine
with 32 grams of hydrogen. The ratio of oxygen in
one compound to another is 32:16 = 2:1.
 Why is this important? It leads to the far more
important understanding that atoms always
combine in whole number ratios to form
chemical compounds (i.e. H2O, never H2.1O0.8).
 This law is true
Dalton’s laws may seem fairly obvious to us, as do many of the
correct principles from these other atomic theories. Let’s see why it’s
not so obvious:
History of the Atom – Modern Theories of Atomic Structure
Thomson (1897): “Plum Pudding Model”
 The model: Atoms are big balls of positive charge with
negative charges (called electrons) embedded in it – imagine a
3-D chocolate chip cookie.
o Because there were no “anode rays”, the positively
charged particles were seen as too big to move from one
place to another.
 First characterization of the electron.
Rutherford (1911): “Gold Foil Experiment”
 How he explained these observations:
o Positively charged alpha particles would pass through the
gold foil most of the time without being deflected.
o The alpha particles would only be deflected when they
passed near the small positively charged nucleus of the
atom (because like charges repel one another)
 Rutherford’s model of the atom:
o Nuclear model: The positive charge is all concentrated
in the middle of the atom, while the negatively charged
electrons float around throughout the rest of the atom’s
empty space.
Our current model of the atom (which we will talk about in far greater
detail shortly) says that the atom contains three types of subatomic
particles that have the following properties:
Particle
Location
Charge
proton (p+)
nucleus
+1
neutron (nº)
nucleus
orbitals surrounding
electron (e-)
nucleus
0
-1
Mass
1 amu
(1.67 x 10-27 kg)
1 amu
~0
(1/1836 amu, or 9.11 x 10-31 kg)
Writing Atomic Symbols:
Atomic symbols have the general format:
Where:
 A is the atomic mass (sometimes called the “mass number”),
which is equal to protons + neutrons. Units of atomic mass
are “amu.”
 Z is the atomic number. This is equal to the number of
protons in all atoms and the number of electrons in neutral
atoms. (Explain how the protons = electrons, making the
charge cancel.) The atomic number determines the
element.
 X is the atomic symbol. If you match the number of protons
to the atomic number on the periodic table, you can figure out
what the symbol is supposed to be.
Isotopes:
For some elements, there can be more than one possible atomic
mass number.
 Why? The neutrons are present to stabilize the protons in the
nucleus so that they don’t repel too much to keep the atom
stable. Different numbers of neutrons can be effective for
making this separation.
Example: Imagine an atom with three protons:
Isotopes: Atoms with the same number of protons (they are the
same element, after all) but different numbers of neutrons.
 There is a common misconception that all isotopes are
dangerous or radioactive.
o This isn’t true. All atoms can be said to be isotopes of
some element, and all elements have isotopes. Only
some isotopes are radioactive, usually those with so
many protons that they can’t hold together no matter how
many neutrons you put in the nucleus.
o All isotopes of uranium and subsequent elements are
radioactive.
If all isotopes have an atomic mass that’s a whole number, why
are all the atomic masses on the periodic table decimals?
 There are many isotopes for each element, and the atomic
mass given on the periodic table is a weighted average of all
their isotope masses.
Unfortunately, in the real world you can’t just count out all of the
atoms and divide the total weight by this number. After all, atoms are
very tiny (~10-10 meters across), making them difficult to manipulate.
How do they find isotopic abundances of elements in the real world?
 Time of flight (TOF) mass spectrometry:
 Sample TOF mass spectrum of magnesium:
The important features of this spectrum:
o m/z represents how far the particles have traveled and is
a direct measure of the mass (m) of each isotope.
o The height of the peak represents the abundance of the
particles. Higher peaks = higher isotopic abundance.
Once you have the isotopic abundances, you can find the average
atomic mass of the elements using the following equation: