TOPIC: Chapter 4 History of the Atom
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DATE____________
TEACHER NOTES
History of the atom
Not the history of atom, but the idea of the atom.
Original idea of at atom came from Ancient Greece (400 B.C.)
Philosophers of the time:
1. Democritus and Leucippus- Greek philosophers who came up
with the concept of an “atom” from the atomos which means
indivisible.
2. Aristotle - Famous philosopher who believed that all
substances were made of 4 elements: Fire – Hot, Air – light,
Earth - cool, heavy, and Water – wet. Blend these in different
proportions to get all substances.
Who was correct: Democritus or Aristotle.
 Greek society was slave based. The famous
did not work with their hands. People did
not experiment. Greeks settled
disagreements by argument. Since Aristotle
was more famous, he won and his ideas
carried through to the middle ages.
Alchemists began experimentataion. They are known for trying to
change lead to gold.
● Late 1700’s - John Dalton an English school teacher who
summarized results of his experiments and those of other scientists.
Teacher- summarized results of his experiments and those of others.
In Dalton’s Atomic Theory, he combined ideas of elements with that
of atoms.
Dalton’s Atomic Theory
1. All matter is made of tiny indivisible particles called atoms.
2. Atoms of the same element are identical, those of different
atoms are different.
3. Atoms of different elements combine in whole number ratios
to form compounds.
4. Chemical reactions involve the rearrangement of atoms. No
new atoms are created nor destroyed.
(Remember the Law and Definite Proportions and Law of
Multiple Proportions.)
Law of Definite Proportions
 Each compound has a specific ratio of
elements.
 It is a ratio by mass.
 Water is always 8 grams of oxygen for each
gram of hydrogen.
Law of Multiple Proportions
 if two elements form more that one
compound, the ratio of the second element
CLASS NOTES
that combines with 1 gram of the first
element in each is a simple whole number.
Atom: The smallest particle of an element that retains the properties
of the element.
Parts of Atoms
 1800’s - Crookes discovers cathode rays (lead to the
TV).
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J. J. Thomson - English physicist. 1897, used a cathode
ray tube to conduct a series of experiments that led to the
discovery of the electron.
 Thomsom’s Model (Plum Pudding Model): A bunch of
positive stuff, with the electrons able to be removed.
 Millikan in 1909 determined the charge of an electron.
 Rutherford, 1911, discovers the nucleus of an atom in his
Gold Foil Experiment.
o Rutherford’s experiment: He believed in the plum
pudding model of the atom. He wanted to see how
big atoms were. He used radioactive Alpha particles
(positively charged pieces given off by uranium). He
shot them at gold foil which was only a few atoms
thick. When the alpha particles hit a florescent
screen, they glowed. (See p. 95).
He expected the alpha particles to pass through
without changing direction very much because he
thought the positive charges were spread out evenly.
Alone they were not enough to stop the alpha
particles.
What he found: the alpha particles were deflected at
different angles. He explained his results by
theorizing that atoms are mostly empty space with a
small dense, positive piece at the center. Therefore,
the positive alpha particles were deflected when they
got close enough.
 In 1919 Rutherford concluded that the nucleus contained
positively charged particles called protons.
 In 1932, Chadwich showed the nucleus also contained
neutrons (no charge).
Subatomic particles
NAME
SYMBOL
CHARGE
Proton
P+
+1
RELATAIVE
MASS
1
Neutron
No
0
1
Electron
e-
-1
0 (1/1840)
Actual
Mass
1.67 X
10-24
1.67 X
10-24
9.11 X
10-28
Modern View
 The atom is mostly empty space.
 Two regions: 1. Nucleus- protons and neutrons. 2. Electron
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cloud- region where you might find an electron.
Density and the Atom: The nucleus has small volume, big
mass (almost all the mass of the atom); thus, big density. The
electron cloud provides most of the volume of an atom, but
virtually no mass.
Size of an atom
 Atoms are small.
 Measured in picometers, 10-12 meters.
 Hydrogen atom, 32 pm radius.
 Nucleus tiny compared to atom.
 If the atom was the size of a stadium, the nucleus would be
the size of a marble.
 Radius of the nucleus near 10-15m.
 Density near 1014 g/cm.
Counting the Pieces
 Atomic Number = number of protons
o # of protons determines kind of atom.
o the same as the number of electrons in the neutral
atom.
 Mass Number = the number of protons + neutrons.
o All the things with mass.
Isotopes
 Dalton was wrong. Atoms of the same element can have
different numbers of neutrons; therefore different mass
numbers. They are called isotopes.
Symbols of isotopes:
 Contain the symbol of the element, the mass number and the
atomic number.
6
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12 C
Naming Isotopes
 Put the mass number after the name of the element.
 carbon- 12
 carbon -14
 uranium-235
Atomic Mass
 How heavy is an atom of oxygen? There are different kinds
of oxygen atoms.
Scientist are more concerned with average atomic mass. It is
based on abundance of each element in nature. Don’t use
grams because the numbers would be too small. Scientist use
units, not grams. The actual mass of a atom in grams is too
small.
Measuring Atomic Mass
 Unit is the Atomic Mass Unit (amu). An amu is one twelfth
the mass of a carbon-12 atom. Each isotope has its own
atomic mass we need the average from percent abundance.
Calculating averages
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If 80% of the rocks were 50 grams and 20% of the rocks
were 60 grams, the average = % as decimal x mass + % as
decimal x mass.
Average mass = .8 x 50 + .2 x 60
Magnesium has three isotopes. 78.99% magnesium 24 with a
mass of 23.9850 amu, 10.00% magnesium 25 with a mass of
24.9858 amu, and the rest magnesium 26 with a mass of
25.9826 amu. What is the atomic mass of magnesium? If not
told otherwise, the mass of the isotope is the mass number in
amu.
o .7899 x 23.9850 + .10 x 24.9858 + .1101 x 25.9826
= 24.305 amu.
Atomic Mass is not a whole number because it is an average
of all the isotopes. It is the decimal numbers on the periodic
table.
Unstable Nuclei and Radioactive Decay
 Nuclear Reactions: involve a change in an atom’s nucleus.
 Radioactivity: the spontaneous emittion of radiation. It occurs
during nuclear reactions. The nucleus changes, so the identity
of the atom changes.
 Radiation: the rays and particles emitted by radioactive
material.
 Radioactive atoms emit radiation because their nuclei are
unstable. Unstable nuclei lose energy to become stable by
emitting radiation through a process called radioactive
decay.
 Types of radiation commonly emitted during radioactive
decay
o Alpha radiation: emits alpha particles. An alpha
particle (α) contains two protons and two neutrons
(2+charge). It = a helium nucleus. Stopped by paper.
226 Ra
222 Rn

+ 42He or
88
86
Radium-226
Radon-222 + alpha particle
o Beta Radiation: emits beta particles. A beta particle
(0-1 β) is a fast moving electron (1- charge). Stopped
by aluminum foil.
14 C
 147N
+ 0-1 β
6
carbon-14
nitrogen-14
beta particle
o Gamma radiation: emits gamma rays (00γ). Gamma
rays are high-energy radiation with no mass and no
electrical charge. Usually accompanies alpha and
beta radiation. Stopped by lead or concrete. Most
dangerous. Gamma radiation alone does not result in
new atoms.
 Nuclear stability: determined by the atom’s ratio of neutrons
to protons. Neutron to proton ration for small atoms is 1:1
and 1.5:1 for large atoms. Decay continues until an atom with
a stable nucleus is formed.
 Nuclear fission: the splitting of a nucleus into fragments
 Nuclear fusion: the combining of atomic nuclei.
 Geiger counter: an instrument used to measure ionizing
radiation.
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Half-life: the amount of time necessary for ½ of the
radioactive material to decay.
o Ex: Bismuth-210 has a half-life of 5 days. If
you start with 100 g, in 5 days 50 g remains,
in 10 days (2 half-lifes) 25 g remain.
NOTE SUMMARY