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Atomic Theory and
Nuclear
Honors Chemistry 6.0
Early Models of the Atom
• Democritus: c. 470-400 BC Greek Philosopher
– matter is composed of tiny, discrete, indivisible
particles called atomos (Greek word meaning
indivisible).
– Ideas based on philosophical speculation
– Theory not accepted due to influence of Aristotle
• An atom is the smallest part of
an element that retains the
chemical properties of that
element. It cannot be broken down
by ordinary means.
Alchemy in the Middle Ages
• Matter composed of 4 elements
– earth, air, fire and water.
• Believed that any substance could be formed:
precious metals to elixirs to cure disease and
prolong life
• Believed in transmutation
– “turning lead into gold” cliché
• Condemned by the Catholic
church
– hidden practices
Contributions of Alchemists
• Produced hydrochloric acid, nitric acid, potash
and sodium carbonate.
• Able to identify the elements arsenic,
antimony, and bismuth.
• Invented and developed laboratory devices and
procedures.
• Laid the foundation for the development of
chemistry as a scientific discipline.
Antoine Laurent Lavoisier
• Law of Conservation of
Matter states that matter is
neither created nor destroyed, it
only changes form.
– 1st to announce that air was made
up of 2 gases – oxygen and azote
(nitrogen)
– Work done on combustion,
oxidation, and gases
• Lavoisier is known as the Father of Chemistry.
• In 1771, at age 28, Lavoisier
married the 13-year-old
Marie-Anne Pierrette Paulze.
• Over time, she proved to be a
scientific colleague to her husband.
• She translated documents and
chemistry books from English.
• She created many sketches and
carved engravings of the laboratory
instruments he used.
• She also edited and published
Lavoisier’s memoirs after his death.
• She hosted parties at which
eminent scientists discussed ideas
and problems related to chemistry.
Engravings of Lavoisier’s Equipment by his wife
Lavoisier was Guillotined May 8, 1794
• An appeal to spare his life so
that he could continue his
experiments was cut short by a
judge saying: "The Republic needs
neither scientists nor chemists; the
course of justice cannot be delayed.”
• One and a half years following
his death, Lavoisier was
exonerated by the French
government.
• When his private belongings
were delivered to his widow, a
brief note was included reading
"To the widow of Lavoisier,
who was falsely convicted."
Joseph Louis Proust: 1799
•Law of Definite Proportion states that
compounds always have the same
elements in the same
proportion by mass.
Ex) the ratio of H:O
in water is always 2:16.
John Dalton:1766-1844
• English schoolteacher
• Some of the original
chemical symbols
from his book:
John Dalton: 1803-1808
Proposed Atomic Theory of Matter:
1. An element is composed of extremely small,
indivisible particles called atoms
2. All atoms of a given element have identical
properties that differ from those of other
elements
3. Atoms cannot be created, destroyed, or
transformed into atoms of other elements
Dalton’s Atomic Theory
(cont.)
4. Compounds are formed when atoms of
different elements combine with one another
in small whole-number ratios
5. In chemical reactions, atoms are combined,
separated, or rearranged
Dalton is credited as being the
Father of the Modern Atomic Theory
Dalton (Billiard Ball) Model
atoms are solid, hard, indivisible spheres
These wooden balls, were the first models made to represent atoms and were used
by John Dalton (1766-1844) to demonstrate atomic theory.
Credit: Science Museum/Science & Society Picture Library
Law of Multiple Proportions
proposed by Dalton
• If 2 or more different compounds are
composed of the same two elements, then the
ratios of the masses of the 2nd element is
always a ratio of small whole numbers
– CO (1.0 g C/1.33 g O)
– CO2 (1.0 g C/2.66g O)
• 2:1 ratio of O in the compounds
– NO (1.0 g N/1.14 g O)
– NO2 (1.0 g N/2.28 g O)
• 2:1 ratio of O in the compounds
Benjamin Franklin: 1706-1790
American statesman/scientist
Ben’s lightning
rod in the
Franklin Institute
In 1752 Benjamin Franklin
•
Experimented with electricity
•
He found that an object can have a positive
or a negative charge.
negative and negative: repel
negative and positive: attract
positive and positive: repel
Michael Faraday (1839)
English scientist
• Hypothesized that atoms
contain electric charge.
• Built 1st electrical motor
• Introduced words such as…
– Ion, electrode, anode and cathode
• A unit of electricity was named after him =
farad
• Static Electricity = electrons move and then
are at rest (grounded)
William Crookes – 1875
English scientist
• Cathode Ray Tube: An
evacuated glass tube with gas at
low pressure
• Electricity is passed through 2
electrodes: cathode (negative)
and anode (positive)
• Light is cast from cathode to
anode (look at the shadow)
• Magnet deflects light – this
proved that particles have
charge and mass.
Crookes’ Conclusion
• Light is composed of negatively charged particles
– Discovered based upon magnet deflection and anode
shadow
Crooke’s
Maltese Cross
You Tube Demo
(CRT)
Applying a
Magnet
Tutor Vista
animation
Wilhelm Roentgen – 1895
German scientist
• In a Crookes tube at
very low gas pressure,
rays of unknown origin
were discovered
• Discovery called X rays
• Glass fluoresced and
the air around the
equipment was ionized
Wilhelm Roentgen’s X ray image
of his wife’s hand
J.J. Thompson: 1897
• English Physicist who
said a cathode ray is
made of electrons, they
have mass (9.1 x 10 g) and
are negatively charged
particles. Thus he
is credited with
“discovering”
electrons.
-28
Cathode Rays and Electrons
•
Cathode Ray Tube: An evacuated glass tube
where a beam of electrons flows from the
cathode (negative electrode) to the anode
(positive electrode.)
J.J. Thomson
• Used Crookes tube (gas discharge tube)
• Applied positive and negative field to a beam of cathode
rays. The deflection was the same for all gases.
• Experimentally proved the existence of the electron (e-)
Cathode Ray
Tube (McGraw Hill)
Discovery of the Electron
11 min
http://www.aip.org/history/electron/jjappara.htm
Thomson
Experimented with hydrogen gas at low pressure
• 2nd beam of particles was moving towards the
cathode, therefore, positive particles
• Deflection of positive ions varied with different gases
• Hydrogen ions had the greatest deflection, therefore,
the smallest positive mass
• Hydrogen ion deflection was smaller than that of the
electron, therefore more massive than an electron
– Hydrogen ion = proton
J.J. Thomson
• Calculated the charge to mass ratio using different cathode metals
and different gases
(e/m= 1.76x108 C/g)
• Measured how much they were deflected by a magnetic field and
how much energy they carried.
• He found that the charge to mass ratio was over a thousand times
higher than that of a hydrogen ion, suggesting either that the
particles were very light or very highly charged.
Credit:Science Museum/Science & Society Picture Library
J.J. Thomson
• Made a bold conclusion:
– Cathode rays were indeed made of particles which he
called “corpuscles," and these corpuscles came from
within the atoms of the electrodes themselves,
meaning the atoms were, in fact, divisible.
• Won a Nobel Prize in Physics in 1906.
J.J. Thomson: 1897
•
•
Thought the atom was made up of these
corpuscles (negative charges) distributed in a sea
of positive charge
Related it to “plum pudding”
Different models of
the plum
pudding model
Robert Millikan:1909
American scientist
1.
2.
Oil drop experiment
Measured voltage to
determine the charge on
one electron =
-1.60 x 10-19 coulomb/e-
3.
Millikan Experiment
12 min
Used Thomson’s charge
to mass ratio to
calculated the mass of
an electron
Mass of 1 electron
= 9.11 x 10-28g
•
•
•
•
•
•
•
•
•
•
•
•
An atomizer sprayed a fine mist of oil droplets into the upper
chamber. Some of these tiny droplets fell through a hole in the
upper floor into the lower chamber of the apparatus.
Next, Millikan applied a charge to the falling drops by
irradiating the bottom chamber with x-rays. This caused the air
to become ionized - meaning the air particles lost electrons.
A part of the oil droplets captured one or more of those extra
electrons and became negatively charged
By attaching a battery to the plates of the lower chamber he
created an electric field between the plates that would act on
the charged oil drops
He adjusted the voltage till the electric field force would just
balance the force of gravity on a drop, and the drop would
hang suspended in mid-air.
Some drops have captured more electrons than others, so they
will require a higher electrical field to stop
Particles that did not capture any of that extra electrons were
not affected by the electrical field and fell to the bottom plate
due to gravity.
When a drop is suspended, its weight m · g is exactly equal to
the electric force applied, the product of the electric field and
the charge q · E.
The values of E (the applied electric field), m (the mass of a
drop which was already calculated by Millikan), and g (the
acceleration due to gravity), are all known values. Unknown
charge on the drop, q
m·g=q·E
Millikan repeated the experiment numerous times varying the
strength of the x-rays ionizing the air so that differing numbers
of electrons would jump onto the oil molecules each time.
He obtained various values for q. The charge q on a drop was
always a multiple of 1.59 x 10-19 Coulombs.
This is less than 1% lower than the value accepted today:
1.602 x 10-19 C
Ernest Rutherford: 1903
• Rutherford studies
under Thomson.
• He discovered 3 types
of natural radiation or
radioactive decay.
α - Alpha Particles
β - Beta Particles
γ - Gamma Rays
high energy X-rays
Rutherford’s Gold Foil Experiment 1909
• This experiment showed the atom has a
small, central positive nucleus and that most
of the atom is empty space.
Gold Foil
Experiment on
You Tube
Rutherford
Video Clip
E drive
You Tube: Discovery of the Nucleus
15 min
Rutherford’s Gold Foil Experiment
Used a narrow beam of  particles to bombard
targets made of thin sheets of gold. Metal foil was
surrounded by a fluorescent screen.
Results:
•most of the  particles passed through the foil
•some were deflected at small angles
•few were deflected at large angles
View of the atoms in the
Gold Foil Experiment
• Rutherford's Gold Foil Experiment
Conclusions:
•atom must contain a very small, dense center of positive charge
•NUCLEUS
•all the positive charge and 99.9% of the mass is in the nucleus
•electrons define the space of an atom
•electrons move at high speeds around the nucleus
•atom does not have uniform density
Gold Foil
Experiment on
You Tube
Rutherford: 1909
• After his Gold Foil
Experiment,
Rutherford modifies
his model of the
atom to contain 2
basic regions: a small
dense positive nucleus
(protons) with
electrons outside.
• Proposed a neutral
part of the nucleus
Neils Bohr: 1913
•
•
•
Thought the atom was like the solar system (planetary
model). Electrons orbit the nucleus with a fixed energy.
Energy Levels - analogous to rungs of a ladder
He wins the Nobel Prize for this model in 1922. It was
eventually shown to be inaccurate and too simplistic.
Henry Moseley: 1913
•
•
•
Worked under Rutherford.
Using CRT’s he bombarded
metals with electrons and
observed the emitted X rays
by the metals
Results: each metal
produced X rays of unique
frequencies or wavelengths
(X ray spectral lines)
Moseley cont.
• Conclusions: He determined that each
element has a unique nuclear charge.
Hence, a different number of protons
(Atomic Number).
• Each atom is electrically neutral and
therefore has an equal number of electrons.
• Killed by a sniper in WW in 1915
James Chadwick: 1932
• Studied under Rutherford.
• 1st isolated a neutron by
bombarding beryllium atoms with
alpha particles
• He determined that the atom also
contained a neutron which had
approximately the same mass as a
proton
– Mass of proton = 1.673x10-24g
– Mass of neutron = 1.675x10-24g
• He proposed that the neutron had a
neutral charge
Chadwick won the Nobel
Prize for his work in 1935.
Wave (electron cloud) Model:
1924 to Present
• Using Quantum Mechanics, the electron can
be found in a probability region.
FUN SONG
The atom through the ages…
The Atom Song
By Michael Ouffutt
To sum it up:
Crash Course on the History of the Atom
Therefore: 
• There are 3 subatomic particles: protons, neutrons
and electrons. These are measured in “atomic
mass units” (amu) as their mass is so small.
Subatomic
Particle
Mass (amu)
Location
Charge
Proton ( p+ )
1.673 x 10-27 kg
(1.0073 amu or 1 amu)
In the nucleus
+
Neutron ( n0 )
1.675 x 10-27 kg
(1.0087 amu or 1 amu)
In the nucleus
0
e- )
9.1x 10-31 kg
(0.0005 amu or 0 amu)
Outside the nucleus
-
Electron (
Atomic Number and Mass Number
• Atomic Number = the number of protons
– Unique to each element
– In a neutral atom, the number of protons
equal the number of electrons.
• Mass Number equal to the total number of
protons + neutrons in the nucleus of an atom.
Ex) carbon-12
Isotopes
Atoms that have the same number of protons
but a different number of neutrons (mass.)
Isotopic Notation
Shorthand way of representing an isotope of an element.
Ex)
37
17
Cl
top number is the mass number (#p + #n)
bottom number is the atomic number (#p)
May also be written: chlorine-37 or Cl-37
The actual average atomic mass for all chlorine isotopes is 35.45 amu
Isotopes of Hydrogen
a. hydrogen (hydrogen – 1)
b. deuterium (hydrogen – 2)
c. tritium
(hydrogen – 3)
Isotope
Carbon-12
Carbon-13
Carbon-14
Protons
6
6
6
Neutrons
6
7
8
1p+
0n0
1p+
1p+
1n0
2n0
Mass
Number
12
13
14
Electrons
1
1
2
1
3
1
H
H
H
Isotopic
Notation
6
12
6
C
6
13
6
C
6
14
6
C
Ions
• Formed when an atom gains or loses an electron
a. Charge = # of protons - # of electrons
Ex) Mg +2 = lost 2 electrons
# of protons: 12 # of electrons: 10 Charge: +2
Positively Charged ion - CATION
Ex) N-3 = gained 3 electrons
# of protons: 7 # of electrons: 10 Charge: -3
Negatively Charged ion - ANION
CATION
“cat”ion
ca+ion
ANION
“ant”ion
Isotope
Mg-25
N-14
Br-79
Protons
12
7
35
Neutrons
13
7
44
Mass
Number
25
14
79
Electrons
Isotopic
Notation
Mg 2
10
25
12
10
14
7
36
79
35
N
3
Br
Charge
+2
-3
1
-1
Atomic Mass:
• The mass of an atom expressed in amu (atomic mass units.)
• One amu is equal to 1/12 the mass of a carbon-12 atom.
Average Atomic Mass:
• The weighted average of all an element’s isotopes.
• Mass Spectrometers are instruments used to measure masses
of isotopes as well as their isotopic abundance.
• This is the number shown in the box on the Periodic Table.
• It is calculated by: (mass1 x %1) + (mass2 x %2) + …
Weighted Average Grade Example:
Straight Class
Weighted Class
Ex) carbon
Ex) hydrogen
93% Tests
90% HW
70% Participation
84.3% Average
x 70% =
x 20% =
x 10% =
Weighted
Average: 90.1%
C-12
C-13
C-14
? Straight
Average
13???
Actual Average
Atomic Mass =
12.011 amu
H-1
H-2
H-3
? Straight
Average
2???
Actual Average
Atomic Mass =
1.0079 amu
Calculation of atomic mass
Magnesium has 3 naturally occurring isotopes:
78.99% Mg-24, 10.00% Mg-25, and
11.01% Mg-26
Calculate the atomic mass of magnesium.
(24 x 0.7899) =
18.9576
=
18.96
+ (25 x 0.1000) =
2.500
=
2.500
+ (26 x 0.1101) =
2.8626
=
2.863
24.323  24.32 amu
Calculation of atomic mass
Magnesium has 3 naturally occurring isotopes:
78.99% is 23.98504 amu
10.00% is 24.98584 amu
11.01% is 25.98259 amu
Calculate the atomic mass of magnesium.
(23.98504 amu x 0.7899)
+ (24.98584 amu x 0.1000)
+ (25.98259 amu x 0.1101)
24.31 amu
Natural Abundance of
Oxygen Isotopes
Isotope
Atomic Mass
(amu)
Natural
Abundance
15.99491
99.759%
16.99913
0.037%
17.99916
0.204%
Average Atomic Mass = 15.9994 amu
Atomic Mass - the mass of an atom, based on a C12 atom, in atomic mass units (amu)
1 amu = 1.66 x 10-24g =
1/12 the mass of a C-12 atom
Example: atomic mass of Na = 23.0 amu
Atoms are too small to count or mass individually. It is easier
count many or mass many.
amu
gram
mole (macroscopic scale)
(atomic scale)
Atomic Mass Units - The Chemistry Journey (3:24)
Mole = amount of substance that contains 6.02 x 1023
particles (abbreviated: mol)
Avogadro’s Number = number of particles in a mole
mole = 6.02 x 1023 particles
Particles can be atoms, ions, molecules, or formula
units
Molar Mass = mass, in grams, per 1 mole of a
substance
units = grams/mole (g/mol)
1 Mole = 6.02x1023 particles of substance
1 Mole = mass (g) of substance from PT
Change the composition of an atom’s nucleus.
Protons & neutrons are called nucleons;
atom is called the nuclide.
1.
2.
3.
Elements may be
converted from one to
another
Particles within the
nucleus are involved.
Tremendous amounts
of energy are released
(million times that of
chemical)
4.
Rate of reaction is not
influenced by external
factors.
1.
2.
3.
4.
No new elements can
be produced
Only electrons
participate
Relatively small
amounts of energy are
absorbed or released
Rate of reaction
depends upon factors
such as temperature
and pressure
Ionizing Radiation is radiation with sufficient energy to
change atoms and molecules into ions (can damage living
tissue).
Nonionizing Radiation is radiation that does not have
sufficient energy to ionize matter.
A. Alpha Decay – spontaneous emission of alpha particle from
the nucleus; from neutron-poor heavy nuclei
226
88
4
Ra  222
Rn

86
2α
204
82
185
79
Pb 
200
80
Au  _________  42 α
181
77 Ir
Hg + 24 α
B. Beta Decay – spontaneous emission of beta particle from
the nucleus; from neutron-rich nuclei
14
6
C 
14
7
N 
228
88
0
1
β
131
53
I
228
Ra 89
Ac + -10 β
β 
0
1
131
54
Xe
Particle
Proton
(Hydrogen nucleus)
Mass (amu) Charge
1.00727647
+1
Neutron
1.00866490
0
Beta Particle
0.0005486
-1
4.00150617
+2
0
0
(electron)
Alpha Particle
(Helium nucleus)
Gamma Ray
(high energy EMR)
Positron
(positively charged
electron)
0.0005486
+1
Symbol
p+ or
n0 or
H
1
0
n
paper
Few cm of lead
e or β
Heavy
clothing/Al foil
He or α
paper
0
1
4
2
1
1
Stopped by
0
-1
4
2
0
0
γ or E
Several cm of
lead
β
Heavy
clothing/Al foil
0
+1
target nucleus
ejected particle
Unstable compound
nucleus
14
7
N  He  [ F ]  O  H
4
2
projectile
18
9
17
8
new isotope
(element)
1
1
9
4
C  n
Be  He  _____
12
6
4
2
238
92
N 
U  _____
14
7
247
99
1
0
Es  5 n
1
0
2
1
n
H  H  He  _____
10
4

Be  N  Mg  _____
3
1
14
7
1
0
4
2
24
12
0
1
How Radon Gas Enters your House
Ways to Remove Radon Gas from Your Home
• External view of a
Radon mitigation system
from a home basement.
• Below is a view of the
fan inside which runs 24
hours a day pulling air
from under the basement
floor.
What is the ½ Life of Strontium-90???
~28 years
How long until no more Strontium-90 remains?
2 days



Used for
determining the age
of previously living
material.
For material up to
25,000 years old,
carbon-14 is used.
For material over
25,000 years old,
potassium-40 is
used.
Half-life problems
1. The half-life of fluorine-18 is 109.8 minutes.
How many hours will it take a 3.60 µg sample
to decay to 0.23 µg?
7.32 hours
2. The half-life for americium-241 is 432 years.
How much of a 50.0 mg sample will remain
after 2590 years?
0.781 mg