History of Atomic Theory
Democritus
– 460-371 B.C.
– ancient Greek philosopher
– believed all matter consisted of
extremely small particles that
could not be divided
– atoms, from Greek word atomos,
means “uncut” or “indivisible”
Aristotle
– believed all matter came from only
four elements—earth, air, fire and
water
Who Was Right?
• Greek society was slave based
• No experiments
– It was all a thought game
• Settled disagreements by argument
• Aristotle was more famous so he won
• His ideas carried through to the middle
ages.
John Dalton (Late 1700’s)
•
•
•
School teacher in England
Based his conclusions on
experimentation and observations.
Combined ideas of elements with that
of atoms
Dalton’s Atomic Theory
1. All elements are composed of
submicroscopic indivisible parts 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 or destroyed.
Parts of Atoms
• Most of Dalton’s theory is accepted today.
• Except the part about atoms being indivisible
J.J. Thomson
and the Cathode Ray Tube
1897
English physicist
Provided the first evidence that atoms are made
of even smaller particles
Description of a cathode ray tube and a short video of
how it works:
http://www.chem.uiuc.edu/clcwebsite/cathode.html
Thomson’s Experiment
-
+
Thomson’s Experiment
-
+
Thomson’s Experiment
-
+
Thomson’s Experiment
-
+
• Passing an electric current makes a beam appear to
move from the negative to the positive end.
Thomson’s Experiment
-
+
• Passing an electric current makes a beam appear to
move from the negative to the positive end.
Thomson’s Experiment
• By adding an electric field
Thompson’s Experiment
• By adding an electric field
Thompson’s Experiment
• By adding an electric field
Thompson’s Experiment
• By adding an electric field he found the
moving particles were negative
Thompson’s Model
• Found the electron
– 1 unit of negative charge
– Mass 1/2000 of hydrogen atom
– Later refined by Millikan to
1/1840
• Concluded that there must be
a positive charge since atom
was neutral
• Atom was like plum pudding
– A bunch of positive stuff, with
electrons able to be removed.
Other Pieces
• Proton – positively charged pieces
1,840 times heavier than the electron
• Neutron – no charge but the same mass
as a proton.
Ernest Rutherford
• Former student of J.J. Thomson
– Believed in plum pudding
1871-1937
• Wanted to find out how big they
are
• Fired positively charged alpha
particles at a piece of gold foil,
which can be made a few atoms
thick
Rutherford’s Experiment
• When alpha particles hit a flourescent
screen it will glow.
• Here’s what it looked like (pg. 90)
• What he expected to see
– Alpha particles should pass through
without change in direction
– Positive charges were spread out evenly.
Alone they were not enough to stop an
alpha particle
• What he got
http://micro.magnet.fsu.edu/electromag/java/rutherford/
• How he explained it
– Atom is mostly empty
– Small dense, positive piece at the center
– Alpha particles are deflected if they get
close enough to positive center
Niels Bohr (1885-1862)
• Electrons have orbits about the
nucleus (planetary theory)
• Electrons could only exist at given
energy levels
• An energy level is where an electron
is likely to be moving
• Energy levels were like steps on a
ladder
– An electron can only be at any given
step at any given time
Modern Atomic Theory
Bohr Model—shows
electrons in orbit
around protons and
neutrons
Quantum-mechanical
model—doesn’t show exact
location of electrons, just
probable place
Structure of the Atom
• There are two regions
– The nucleus
• Protons and neutrons
• Positive charge
• Almost all of the mass
– Electron cloud
• Most of the volume of an atom
• Region where electron can be
found
Subatomic particles
Counting the pieces
• Atomic number = number of protons
– Same as the number of electrons in a
neutral atom
• Mass number = the number of protons
+ neutrons
Atomic Mass Unit AMU
• Mass of a proton = 1.67 x 10 -27g
– A pretty inconvenient number
– New unit referenced to mass of an isotope
of carbon: carbon -12
– Carbon-12 has 6 protons and 6 neutrons
• Has a mass of 12.00000 amu – an atomic
mass unit
– Therefore 1 proton and 1 neutron has a
mass of 1 amu.
So why not whole numbers for
atomic masses in periodic table?
• Reported numbers are average atomic
mass units, reflecting the abundance of
isotopes for any given number.
• In nature most elements occur as a
mixture of two or more isotopes
Isotopes
• Atoms of the same element can have
different numbers of neutrons
• Different mass numbers
• Called isotopes