Download History of Atomic Theory

History of Atomic Theory
Unit 3 Lesson 1
Topics
• Famous scientists, their experiments, and how
the atomic model has evolved over time as a
result of their work
– Democritus, Dalton, Thomson, Rutherford,
Chadwick, Bohr
– Cathode Ray Tube exp., Gold Foil exp., Flame
Tests
– Spheres ->Plum Pudding ->Nuclear Atom->
Planetary model
Learning Goals
• To describe famous scientists, their
experiments, and how the atomic model has
evolved over time as a result of their work
• To describe how the planetary model and the
concept of quantized energy levels explains
the atomic emission spectra and phenomena
such as fireworks and colored flames
Atomic Theory Timeline
Democritus: 460-370 B.C.
• 1. There are basic elements from which all
matter is made
• 2. Everything is made of small atoms moving in
a void
• 3. Some atoms are round, pointy, oily, have
hooks, etc. to account for their properties
• 4. Ideas rejected by leading
philosophers
John Dalton: 1766-1844
• 1. Each element is composed of extremely
small indivisible particles called atoms
• 2. All the atoms of a given element are
identical, but different from those of any
other element
• 3. Atoms are neither created nor destroyed in
any chemical reaction
• 4. A given compound always has the
same relative numbers and kinds of
Atoms. (Ex water is always H2O)
John Dalton: 1766-1844
• Where Dalton was wrong
– 1. We can divide the atom
• Subatomic particles include protons, neutrons,
& electrons
• E = mc2  Atomic Bomb
– 2. Not all atoms of a given element are
identical
• Isotopes
J.J. Thomson: 1856-1940
• 1.Discovered electron in 1897
– Cathode Ray Tube Experiment
– Electron has (-) charge: “e-”
– Mass of e- = 9.11 x 10-28g
• 1/1840 the mass of a proton
– Protons discovered in 1886 by Goldstein
• Proton has (+) charge: “p+”
• Mass of p+ = 1.67 x 10-24g
J.J. Thomson: 1856-1940
• Cathode Ray Tube Experiment
• The beam Is attracted to (+) plate and repelled
by (-) plate so it must be made of (-) charged
particles!
J.J. Thomson: 1856-1940
• Thomson knew the atom as a whole was
neutral so there must be (+) charged particles
also.
J.J. Thomson: 1856-1940
• 2. Developed Plum Pudding model in 1904
– Atom is a sphere of (+) charge with (-)electrons
randomly scattered throughout
– Modern version: “Choc. Chip Cookie” ?
J.J. Thomson: 1856-1940
• 3. Discovered isotopes in 1913
– Different “versions” of the same atom. Atoms
of the same element, but with different
masses.
• Nucleus and neutrons had not been
discovered yet so he didn’t fully understand
why isotopes existed
Ernest Rutherford: 1871-1937
• 1. Discovered Nucleus w/“Gold Foil
Experiment” in 1909
– Bombarded a piece of gold foil (gold atoms)
w/large, (+) charged alpha particles.
– Most passed right through but a few bounced
back or were deflected at an angle
Ernest Rutherford: 1871-1937
• 2. Developed Nuclear Atom Theory in 1910
– 1. The atom is mostly empty space
– 2. In the center of the atom is a tiny nucleus,
which contains most of the atom’s mass
– 3. The nucleus has a (+) charge
– 4. Electrons are in afixed orbit in the empty
space around the nucleus
Ernest Rutherford: 1871-1937
James Chadwick:
1891-1974
• Discovered the neutron in 1932
– Neutron has no charge: “n0” – Mass of n0: 1.67 x
10-24g
• Almost the same mass as a proton (both are much
larger than the e-)
• Described the nucleus and explained isotopes
discovered by JJ Thomson in 1913
– Isotopes of the same element have different
masses (same # protons but different #s of
neutrons)
James Chadwick: 1891-1974
Niels Bohr: 1885-1962
• Developed Planetary Model in 1913
• Nucleus is surrounded by electrons that are
orbiting in definite paths
“like planets orbit the sun”
Niels Bohr: 1885-1962
• 2. Electrons exist in discrete “quantized”
energy levels.
• • Quantized = the opposite of continuous,
goes from one level to the next with nothing
in between
Niels Bohr: 1885-1962
• The further you get from the nucleus, the
higher the energy level
• Electrons can absorb a photon (packet of
energy) to move from the ground state
(lowest energy level) to the excited state (one
or more levels higher)
• When they fall back down, they emit energy in
the form of light
– “electromagnetic radiation”
Niels Bohr: 1885-1962
• The electromagnetic radiation emitted has a
certain frequency and wavelength (light
travels as a wave), which corresponds to a
certain color on the visible spectrum
Niels Bohr: 1885-1962
• Every line of color on an atom’s emission line
spectra represents its electrons dropping back
down from a higher to a lower energy level
after being in an excited state
Niels Bohr: 1885-1962
• Each element's emission spectrum is unique
– This is the science behind fireworks and
colored flames
– You can see line spectra with special
equipment