Download Atoms, Molecules and Ions History

Atoms, Molecules and Ions
History
• Greeks
• Democritus 460-370 BC “atomos”
• Aristotle- elements.
• Alchemy
• 1660 - Robert Boyle- experimental definition
of element.
• Lavoisier (1734-1794)- Father of modern
chemistry. He wrote the book -1789.
Laws
• Conservation of Mass
• Law of Definite Proportion- compounds
have a constant composition.
• Multiple Proportions- When two elements
form more than one compound, the ratios of
the masses of the second element that
combine with one gram of the first can be
reduced to small whole numbers.
Multiple What???
• Water has 8 g of oxygen per g of hydrogen.
• Hydrogen peroxide has 16 g of oxygen per
g of hydrogen.
• 16/8 = 2/1
• Small whole number ratios.
Dalton’s Atomic Theory
1803-1807
1) Elements are made up of atoms
2) Atoms of each element are identical. Atoms
of different elements are different.
3) Compounds are formed when atoms
combine. Each compound has a specific
number and kinds of atom.
4) Chemical reactions are rearrangement of
atoms. Atoms are not created or destroyed.
A Helpful Observation
• Gay-Lussac- under the same conditions of
temperature and pressure, compounds
always react in whole number ratios by
volume.
• Avogadro- interpreted that to mean that at
the same temperature and pressure, equal
volumes of gas contain the same number of
particles.
• (called Avogadro’s Hypothesis)
Thomson’s Experiment (1897)
• Used the Cathode Ray Tube(CRT) to
discover the electron.
Thomson’s Experiment
Passing an electric current makes a beam
appear to move from the negative (cathode)
to the positive end (anode).
Proving Electrons had Mass
• Thomson devised a cathode ray tube
with a paddle wheel built inside. When
the high voltage electricity was turned
on the paddle wheel began to rotate and
move away from the cathode and
towards the anode.
Determining Charge
• Thomson concluded from this evidence
and from his previous experiments that
tiny particles were being emitted from the
atoms of the cathode. These tiny particles
were negatively charged. He called these
particles "electrons."
Thomson’s Model
of the Atom
• Given these experimental results,
Thomson proposed in 1897 what has
since been referred to as the "plum
pudding" model of the atom. This model
depicts the atom as a diffuse cloud of
positive charge with the negative
electrons embedded randomly in it, like
plums in pudding.
Millikan (1909) - Mass of the
Electron “(Oil Drop Experiment)”
• Using this information he calculated the mass
of the electron as 9.11 X 10-28 grams.
Radiation
• In 1896, the French scientist, Henri
Becquerel, found that a piece of a mineral
containing uranium could produce its
image on a photographic plate in the
absence of light. He attributed this
phenomenon to a spontaneous emission
of something which he called "radiation"
which originated from the uranium.
Radioactivity
• Discovered by accident
• Bequerel
• Marie and Pierre Curie isolated the
radioactive components at Bequerel’s
suggestion
• Three types
• –alpha- helium nucleus (+2 charge, large
mass)
•
__beta-
high speed electron
• –gamma- high energy light
Rutherford (1911)
• A fluorescent screen would detect
radiation by flashing whenever it was
struck by the radiation.
• Rutherford observed that the radiation
was diffracted into three beams by the
charged plates.
Rutherford “Gold Leaf Experiment”
• Rutherford set up the apparatus which
would bombard thin gold foil with alpha
particles. These particles would then be
detected by the fluorescent screen.
He anticipated that all of the alpha
particle detection would occur on the
screen directly behind the foil.
Rutherford got surprisingly different results
• "It was about as credible as if you had
fired a 15-inch shell at a piece of tissue
paper and it came back and hit you!"
Atomic Model Revised
• Rutherford suggested that the atom was
mostly empty space with a highly charged
center. Most of the particles pass through
the atom undisturbed, but a few get too
close to the center and are deflected.
"Planetary Model."
• To account for these results Rutherford
proposed a new model of the atom in
1913. This model had the following
characteristics:
The atom is mostly empty space with the
majority of its mass concentrated in the
center of the atom which he called the
"nucleus."
Mass Number and Atomic Mass
• By the early 1930's the major subatomic
particles had been discovered and their
physical properties had been described.
• James Chadwick (1932) - discovers
Neutrons.
Particles
Charge
Proton
Mass (grams) Relative Mass (amu)
1.67262 X 10-27
1.007
Electron 9.10939 X 10-31 5.486 X 10-4 (~ 0)
Neutron
1.67493 X 10-27
1.009
Relative
+1
-1
0
Atomic Mass Unit
• amu - the unit that we use to measure
atoms.
• 1 amu ~ mass of one mole of the following:
~hydrogen ~ 1 proton ~ 1 neutron
• Actual masses:
• Particle
Charge
Mass (amu)
• Proton
positive(+1)
1.0073
• Neutron
none(neutral) 1.0087
• Electron
negative (-1)
5.486 x 10-4
• Hydrogen
none (neutral) 1.0079
Measuring Atoms
o
Angstrom (A) - a convenient non-SI unit of
length used to express atomic dimensions.
• 1 Angstrom = 1 x 10-10 meters
-10 m
• most atoms have diameters between
1
x
10
o
and 5 x 10 -10 m, or between 1 - 5A.
Isotopes
• All atoms have the same number of protons.
• The number of neutrons may vary for a
given element.
• Isotope - atoms of a given element that
differ in the number of neutrons and
consequently the mass.
• May be written as the symbol 12C or simply
6
carbon-12 as opposed to the isotope 14C or
6
carbon-14.
Atomic Numbers and Mass Numbers
• Atomic Number - is the number of protons, which is
shown as the subscript. (it is also the number of
electrons in a neutral atom.)
• Mass Number - is the total number of protons plus
neutrons in the atom. (which represent essentially
all the mass of the atom.)
12C
• ex.
6
• has 6 Protons, 6 Electrons and 6 Neutrons
• Number of Neutrons = Mass # - Atomic #
6
=
12 6
Nuclides
• Nuclide - an atom of a specific isotope
14C or carbon-14
• ex.
6
• 6 protons, 6 electrons and 8 neutrons
• # of neutrons = atomic mass - atomic number
•
8
=
14
6