Chapter 3
Atomic Theory
Today’s Objectives
• Understand the basics of Dalton’s Atomic Theory,
and how it relates to the study of chemistry; be
aware of how it differs from the currently accepted
modern Atomic Theory.
• Know the basic details of how each subatomic
particle was discovered, and what information was
determined about protons, neutrons, and electrons
in these experiments
Early Theories
• 4 elements
Democritus
(460 - 370 BC)
• Greek philosopher
• Atomos –
indivisible
particles
• Atoms are the
smallest particle
that retains the
chemical identity
Jabir Ibn Haiyan
(700? - 803 AD)
• Father of
Chemistry
• Practiced
Alchemy
• Discovered
metals
Lavoisier
(1743 - 1794)
• Law of
Conservation
of Matter
Joseph Louis Proust
(1754-1826)
• Law of
Constant
Composition
Dalton
(1766 - 1844)
• Atomic Theory
Atomic Theory
• Elements made of atoms
• Atoms are identical of a given
type of element
• Atoms neither created nor
destroyed
• Compounds have fixed ratio of
atoms
Ben Franklin
(1706-1790)
• Two types of
charge positive
(+) and
negative (-)
Michael Faraday
(1791-1867)
• Atoms are
related to
electricity
J.J. Thomson
(1856 - 1940)
• Cathode Ray
Tube (CRT)
stream of
electrons
• Plum Pudding
Model
cathode ray
tube
Robert Millikan
(1838-1953)
• Determined
charge & mass
of electron
Becquerel
(1852 - 1908)
• Uranium
exposes film
Marie (1867-1934) & Pierre
Curie (1859-1906)
• Discovered
radioactivity
elements
• Radioactive
decay
Rutherford
(1871 - 1937)
• Discovered
radioactivity
particles
• Discovered
Nucleus
• Solar system
model of atom
Discovery of particles
gold foil experiment
Niels Bohr
(1885 – 1962)
• Electrons do not
orbit like
planets
• Described shells
or energy levels
• Quantum theory
H.G.J. Moseley
(1887 - 1915)
• Discovered
protons (+) in
the nucleus
• Rearranged
periodic table
Sir James Chadwick
(1891-1974)
• Discovered
neutrons (0) in
the nucleus
Today’s Objectives
• Review History of the atom
– Dalton’s Theory
– History and discovery of each subatomic particle
• Discovery Learning:
– Know the name, location, charge, and relative mass of each of the
subatomic particles in an atom
– Know that the atomic number is the number of protons in the
nucleus of an atom, and is unique to each element.
– Understand that isotopes are atoms of the same element that differ
in the number of neutrons in the nucleus, and therefore differ in
mass.
– Know the mass number is, and be able to use it to correctly
designate isotopes using both hyphen notation and nuclear
symbols.
Protons
•
•
•
•
•
Make up the nucleus
Charge +1.602 x 10 -19C
Mass = 1.673 x 10 -24g
Charge +1
Mass = 1 amu
Neutrons
•
•
•
•
Make up the nucleus
Charge 0
Mass = 1.675 x 10 -24g
Mass = 1 amu
Quarks, Quarks, Quarks
(1950s – present)
• 6 quarks have
been
discovered
that make up
protons and
neutrons
Electrons
•
•
•
•
•
Occur in electron Clouds
Charge -1.602 x 10 -19C
Mass = 9.109 x 10 -28g
Charge = -1
Mass = 0 amu
• Atoms are small but nuclei are
smaller
• Diameter of a penny has 810
million copper atoms
Atomic Number
• Number of protons in an atom
• Electrically neutral atoms have the
same number of electrons as
protons
• Ions are formed by gaining or
losing electrons
Atomic Mass
• The mass of the nucleus, number of protons
and neutrons.
• It is an average of all the isotopes masses
Isotopes
• Same number of Protons but
different numbers of neutrons
• Mass number is the sum of the
protons and the neutrons
• Isotopes have the same chemical
properties
• Violates Dalton’s atomic theory
Masses of Atoms
•
•
•
•
•
•
1 amu = 1/12 mass of a 12C atom
99% Carbon 12C
1% Carbon 13C
Average atomic mass of C is 12.01 amu
Mass number is for one atom
Listed as a decimal on the periodic
table
Nuclear Symbol
• Hyphen notation
• Li - 7
Today’s Objectives
• Review History of the atom
• Review the following
– Know the name, location, charge, and relative mass of each of the
subatomic particles in an atom
– Know that the atomic number is the number of protons in the nucleus of
an atom, and is unique to each element.
– Understand that isotopes are atoms of the same element that differ in the
number of neutrons in the nucleus, and therefore differ in mass.
– Know the mass number is, and be able to use it to correctly designate
isotopes using both hyphen notation and nuclear symbols
• Know the difference between a neutral atom and an ion is the number
of electrons, the charge of the ion is the difference between the number
of protons and neutrons
• For any given atom or ion, be able to determine the number of each
subatomic particles.
IONS
• Same number of protons but a different
number of electrons
• Charged particles because the number of
electrons (-) is either greater than or less
than the number of protons (+)
• H+
• Ions are determined by there place on the
periodic table.
Practice
• Determine the number of subatomic
particles for the following:
–
–
–
–
He
C -14
Na+
O-2
Nuclear Reactions
• Nuclear reactions involve the nucleus
of the atom
• Radioactivity is the spontaneous
emission of radiation from an atom
• Nuclear reactions change elements
involved
Alpha Particle
• Alpha particle
– Helium nucleus
with no electrons
– Will bounce off
of paper and skin
– +2 charge
Beta Particle
• Beta particle
– High energy electron
– Come from the decay
of a neutrons
– Will penetrate skin
– Blocked by
aluminum and
Plexiglass
– -1 charge
Gamma Radiation
• Gamma Rays
– High energy wave
– No charge
– No mass
– Penetrates skin, damages cells and mutates
DNA
– Blocked by lead
Nuclear Stability
• Most elements have a
stable nucleus
• A strong nuclear force
holds protons and
neutrons together
• Neutrons act as the
“glue” holding the
protons together
Nuclear Equations
• Scientists use a nuclear equation
when describing radioactive decay
• The mass number and atomic
number must add up to be the same
on both sides of the equation
Beta Decay
• Beta decay results in an increase in
the atomic number
Practice
• Write the nuclear equation of the
alpha decay of Radon – 226
• Write the nuclear equation of the
alpha decay of Gold - 185
Practice
• Write the nuclear equation of the
beta decay of Iodine - 131
• Write the nuclear equation of the
beta decay of Sodium - 24
Chapter 24
Applications of Nuclear
Chemistry
Half Life
• Radioisotopes are radioactive
isotopes of elements (not all isotopes
are radioactive)
• A half-life is the amount of time it
takes for one half of a sample to
decay.
• http://lectureonline.cl.msu.edu/~mmp/applist/de
cay/decay.htm
Beta Decay of Phosphorous - 32
Radiocarbon Dating
• Carbon - 14 undergoes beta decay
• Half life of 5,730 years
• Used to approximate ages 100 –
30,000 years
• Other radioisotopes are used to
measure longer periods of time
Parent
Daughter
Half Change in...
Carbon-14
Nitrogen-14
5730 years
Uranium-235
Lead-207
704 million years
Uranium-238
Lead-206
4,470 million years
Potassium-40
Argon-40
1,280 million years
Thorium-232
Lead-208
14,010 million years
Rubidium-87
Strontium-87
48,800 million years
Nuclear Bombardment
• Nuclear scientists make nuclei
unstable by being bombarded with
particles
• Also known as particle
accelerators or “atom smashers”
Radiation
• SI units are in Curies (Ci)
• One Curies is amount of nuclear
disintegrations per second from
one gram of radium
• Also measured in rem (Roentgen
equivalent for man
• Over 1000 rem is fatal
• Detected by a Geiger counter
Nuclear Power
• Nuclear Reactors use fission of
Uranium-235 as source of energy
• A large nucleus is split into two
smaller nuclei
• A small amount of mass is converted
to a tremendous amount of energy
• ~1 lb Uranium 235 = 1 million gallons
of gasoline
• http://people.howstuffworks.com/nuclearpower2.htm
Nuclear Fusion
• 2 atomic nuclei
fuse releasing a
tremendous
amount of
energy
Nuclear Weapons
• Source of
energy is
Plutonium or
Hydrogen
• Can be fusion
or fission
Gun-triggered fission bomb
(Little Boy - Hiroshima),
Implosion-triggered fission
bomb (Fat Man - Nagasaki),
http://people.howstuffworks.com/nuclear-bomb5.htm