Atomic Structure – History and Introduction
Ancient Greek philosophers had 2 opposing views of the fundamental nature of matter.
Democritus (circa 400 BCE)
Matter is made up of discrete, indivisible particles called atoms.
Aristotle (circa 300 BCE)
Matter is continuous and can be infinitely divided.
Neither view was based on experimentation, but in a series of debates Aristotle’s view
was favored and influenced scientific thinking until the 1700’s.
In the 1700’s two experimentally based concepts became widely accepted:
1) An element is a substance that cannot be broken down into simpler substances by
ordinary physical or chemical means.
2) Elements combine via chemical reactions to form compounds that have different
physical and chemical properties than the original elements.
These concepts started chemists questioning the Aristolian view of matter as being
continuous.
The other major trend in chemistry that was taking place in the 1700’s was an emphasis
on Quantitative Analysis. This trend was brought about by improvements in
technology (better balances, thermometers, etc.) allowing for greater accuracy and
precision. The major proponent of this move toward quantitative chemistry was
Antoine Lavoisier.
2 basic laws grew out of this movement:
1) Law of Conservation of Mass (Matter) – Antoine Lavoisier, 1789
Matter is neither created nor destroyed in ordinary chemical or physical changes.
Mass before = Mass after
Demo: Reaction on a balance (BaCl2 + Pb(NO3)2 Ba(NO3)2 + PbCl2)
2) Law of Definite Proportions (Constant Composition) – Joseph Louis Proust, 1799
A compound always contains the same elements, in the same proportions by
mass, regardless of the size or source of the sample.
i.e.:
Water - H2O has a 2:1 particle or molar ratio
H = 2(1.00794) = 2.01588
O = 15.9994
Total mass = 18.01528
The H:O mass ratio is approximately 2:16 or 1:8
John Dalton, 1803 explained these various findings by proposing that elements are
made up of atoms that combine to form compounds. His proposal included:
1) All matter is composed of small particles called atoms.
2) All atoms of a given element are identical in size, mass and properties; while
atoms of different elements differ in these characteristics.
3) Atoms are indivisible (cannot be split, created of destroyed in ordinary
chemical or physical changes). {Modification for modern atomic theory}
4) In chemical reactions atoms can be combined, separated or rearranged.
5) Atoms combine in simple whole-number ratios to form chemical compounds.
As an out-growth of these five ideas Dalton also stated a law called the Law of Multiple
Proportions, which states:
If 2 elements combine to form more than one compound, with one element being
present in a fixed amount, the different masses of the other element will be in
small, whole number ratios.
Example 1:
Hydrogen and Oxygen can combine to form Water (H2O) or Hydrogen peroxide (H2O2).
We already found the mass ratio for Water to be 1:8
The mass ratio for Hydrogen peroxide is:
Hydrogen = 2(1.00794) = 2.01588
Oxygen = 2(15.9994) = 31.9988
These rounds off to be a 2:32 or 1:16 ratio. Since Hydrogen in both ratios is 1
then the ratios between the Oxygens is 8:16 or 1:2.
Dalton pictured the atom as a solid sphere.
(Slide 1)
During the 1700’s and 1800’s experiments were being done by numerous scientists,
including Benjamin Franklin, on the nature of electricity. Franklin proposed that objects
could have two possible charges, which he named positive and negative. He proposed
that these charges interacted in a simple manner, similar charges repel while opposite
charges attract. In 1839 Michael Faraday suggested that chemical properties and atomic
structure are related to these electrical attractions and repulsions.
Joseph John Thomson, 1890’s performed a series of experiments with the cathode ray
tube including: (Demos)
1) Deflection of the path by electrical/magnetic field – showed that the cathode
ray had a negative charge.
(Slides 2 – 4)
2) Maltese cross shadow demonstrated that the Cathode ray was made up of
particles he called electrons.
(Slide 5)
3) Paddle wheel further demonstrated that the Cathode ray was made up of
particles.
(Slide 6)
Thomson was able to determine the mass to charge ratio for the electron and show that
there was a large relative charge in a very small mass.
Robert A. Millikan, 1909 (Slide 7) performed an experiment in which he sprayed
droplets of oil into a chamber with charged plates. By manipulating the charges he was
able to calculate the charge on a drop of oil. Knowing the mass of the oil drop and
using the mass:charge ratio found by Thomson, Millikan was able to calculate the mass
of the electron to be 9.109 x 10-28 g. This is around 1/1837 the mass of the simplest
Hydrogen atom.
Thomson’s and Millikan’s experiments altered the theory of the atom in two ways:
1) Since the atom is neutral there must be positive charges in the atom to balance
the negative charges of the electron. Thomson imagined the atom as having a
positive gel-like matrix with electrons embedded in it. This is called the Plum
Pudding Model of the atom.
(Slide 8)
2) Since electrons have so little mass, compared to the mass of the atom, there
must be other more massive particles present in the atom.
Ernest Rutherford,1903 determined that radiation had 3 components:
(Slide 9)
Alpha particles which have a positive charge
Beta particles with a negative charge
Gamma radiation which is a form of energy similar to light, no charge
Sir Ernest Rutherford, 1911 (Slide 10) with the assistance of Hans Geiger and Ernest
Marsden did an experiment that radically changed our concept of atomic structure.
They firing a stream of alpha particles, (+2 charge with a mass 4x that of a hydrogen
atom) at a thin sheet of gold foil.
He discovered most of his
positively charged “bullets”
passed right through the gold, a
few were deflected and a very
small number bounced straight
back.
He realized the atom’s positive charged particles
were contained in the nucleus. The negatively
charged electrons were orbiting around the nucleus.
(Slide 11)
The particles making up the nucleus are not
discovered until much later.
Niels Bohr, 1913 (Slide 12) based on his studies of the excitation of the Hydrogen atom.
Bohr determined that electrons move in definite orbits around the nucleus, much like planets
orbit around the sun. These orbits, or energy levels, are located at certain distances from the
nucleus.
Each orbit has a limited number of electrons that can occupy that orbit. These limits are:
1st orbit = 2 e2nd orbit = 8 e-
3rd orbit = 18 e4th orbit = 32 e5th orbit = 32 e6th orbit = 18 e7th orbit = 8 eWith a special additional limit that no outer most orbit can hold more than 8 e-, this is
called the Octet Rule.
Rutherford, 1919 discovered the proton (with a +1 charge, and a mass around the
same as an atom of Hydrogen.
Chadwick, 1932 discovered the neutron (no charge and a mass equal to that of a
proton).
Bohr – Rutherford Model, (Slide 13)