Unit 3 – History of Atomic Theory
Democritus of Abdera (c. 460 – c. 370 BCE)
A greek philosopher that believed all matter was composed of discrete, indivisible
particles called στομωσ (atomos). The more accomplished philosophers (Aristotle and
Plato) did not accept this theory. They believed all matter was made of four elements:
earth, air, water and fire.
Antoine Lavoisier (1743-1794)
Following extensive studying of combustion and the discovery of carbon, dioxide,
oxygen, nitrogen and hydrogen gases, Lavoisier regarded measurement as the essential
operation of chemistry. Through careful experiments, he weighed the reactants and
products of various reactions and suggested the LAW OF CONSERVATION OF MASS
Joseph Proust (1754-1826)
Proust furthered Lavoisier’s experiments by weighing the ratios of reactants in
specific chemical compounds. He found that for specific compounds the ratio of their
combining was constant. This led to the LAW OF DEFINITE PROPORTION.
John Dalton (1766-1826)
Dalton developed the first atomic model of the atom that offered explanations for
the simple laws that govern chemistry. solid sphere model
His theory has three main postulates.
1. An element is composed of tiny particles called atoms. All atoms of a given
element show the same chemical properties. *This is not true today because of
the discovery of neutrons and isotopes.
2. Atoms of different elements have different properties. In an ordinary chemical
reaction, no atom of any element disappears or is changed into an atom of
another element.
3. Compounds are formed when atoms of two or more elements combine. In a
given compound, the relative numbers of atoms of each kind are definite and
constant. In general, these relative numbers can be expressed as integers or
simple fractions. This led to the LAW OF CONSTANT COMPOSITION.
Henri Becquerel (1788-1878)
Becqueral accidentally found that a piece of a mineral containing uranium could
produce an image on a piece of photographic paper in the absence of light. He attributed
this phenomenon to a spontaneous emission of radiation by the uranium, which he called
radioactivity. Studies demonstrated three types of radioactive emission: gamma (γ) rays,
beta (β) particles and alpha particles (α) particles. A gamma ray is essentially high energy
“light” (no mass and no charge); a beta particle is a high speed electron (essentially no
mass and a negative charge); and an alpha particle is essentially a helium nucleus (heavy
mass = 4 and 2+ charge). More models of radioactivity have been discovered but the alpha
particles were used in early crucial experiments.
J J Thomson (1856 – 1940)
Thomson did the first important experiments that led to the understanding of the
composition of the atom. He studied electrical discharges in partially evacuated tubes
called cathode ray tubes.
Thomson found that when high voltage was applied to the tube a “ray” (he called it the
cathode ray) was produced. Due to the fact that the ray emanated from the negative pole of
the electric field the ray was composed of negatively charged particles. This led to the
discovery of electrons. Also from this discovery he calculated the charge-to-mass ratio of
the electron. He reasoned that all atoms contain electrons and that to be electrically
neutral, atoms must also contain positive charges. Thomson postulated that an atom
contained a diffuse cloud of positive charge with negative electrons embedded in it. This is
often called the plum pudding or raisin bun model of the atom because the electrons are
like raisins dispersed in the dough (the positive charge).
plum pudding model
spherical cloud of
positive charge
electrons
s
Max Planck (1858-1947)
Planck studied the radiation profiles emitted by solid bodies that were heated to
incandescence. From these experiments, he determined that the results could not be
explained by the typical (Newtonian) physics of the time. The loss or gain of energy could
only be obtained by whole number multiples of the quantity, hν, where h is Planck’s
constant (6.626 x 10-34 J·s/particle) and ν (nu) is the frequency of the electromagnetic
radiation. A small “packet” of energy was called a quantum (plural, quanta).
Robert Millikan (1868-1953)
Millikan performed experiments with charged oil drops. These experiments allowed him
to determine the magnitude of the electron charge. Using this value and the charge-tomass ratio determined by Thomson, Millikan was able to determine the mass of an
electron.
Oil drop experiment
Ernest Rutherford (1871-1937)
Rutherford, who also performed many experiments exploring radioactivity, carried
out an experiment to test Thomson’s model of the atom. The experiment directed alpha
particles at a thin sheet of gold foil. He reasoned that if Thomson’s model were accurate,
the massive alpha particles would crash through the foil with little deflection. Although
most of the particles passed straight through, many of the particles were deflected at large
angles and a very few were reflected back at him. This led to three conclusions:
1. The atom is mostly empty space. Most alpha particles went straight through.
2. The atom has a positively charged center or nucleus. Due to the fact that some
particles came close to the center and the deflection path was more similar to an
electric field of like charges.
3. The nucleus is the concentration of the atom’s mass. The very few that were
deflected were on a collision course with a much more massive center of charge.
Gold Foil Experiment
If Thomson was correct with his model, then a completely different outcome would have
occurred. This Gold-Foil Experiment created the nuclear model of the atom, with the
positively charged nucleus at the center.
James Chadwick (1891-1974)
Capitalizing on the theories proposed by Rutherford and others, Chadwick produced
experimental evidence of the existence of the atomic particle that had a mass nearly equal
to that of a proton but no charge. This was a result of alpha particle bombardments of
beryllium, lithium and boron. He named this particle the neutron according to the
equation: 9Be + α → 12C + 1n
Louis de Broglie (1892-1987)
Utilizing Albert Einstein’s special theory of relativity (energy has mass) and the
incorporations of Planck’s equation, de Broglie supported the view that light could not only
be quantized but that it also could possess wave properties as all matter can. This is called
the wave-particle duality of light. It depends upon the nature of the quantity you are
investigating. Light, which was previously thought to be purely wavelike, was found to
have particle (masslike properties) and the opposite is also true. The relationship is called
de Broglie’s equation λ = h/mv, where λ is the wavelength (in meters), h is Planck’s
constant, m is the mass of the particle (in kilograms), and v is the velocity of the particle (in
meters/second).
Niels Bohr (1885-1962)
As Rutherford’s graduate assistant, Bohr held the planetary model of the atom and
incorporated the work of Planck and de Broglie to propose that the electrons must be
balanced by the attraction for the nucleus to resist flying off the atom. However a
constantly accelerating particle (like the electron) should lose energy and then eventually
fall into the nucleus. Due to this apparent contradiction, Bohr applied a theory to the
hydrogen atom that claimed that energy could only occur in specified increments. These
increments were called energy levels and could also have only whole number values.
Electrons in a normal state (ground state) could gain energy and be promoted to a higher
level (excited state) but eventually that energy could not be sustained and would be
emitted as electromagnetic radiation (visible light in some cases). The specific lines of
radiation emitted were consistent with observed emission spectra. This lead to the
quantized planetary model of the (hydrogen) atom and Bohr’s equation: E = -B/n2, where E
is energy of radiation, B is Bohr’s constant, and N is the energy level. Also energy changes
can be calculated for the hydrogen atom: ΔE = Ehi - Elo = -1312 kJ/mol (1/nhi2 - 1/nlo2)
where “hi” represents the excited level and “lo” represents the ground state. Bohr’s
findings were in agreement with 0.01% of observed results. However, when investigating
helium or higher “multi-electron” atoms the percent error grew exponentially.
By the mid – 1920’s it became obvious that Bohr’s model could not be made to work for
these higher atoms. Three physicists, Werner Heisenberg, Louis de Broglie and Erwin
Schrödinger (1887 - 1961)
proposed a new branch of physics to help explain the
contradictions. This new branch of physics is called quantum mechanics. Schrödinger gave
greater emphasis to the wave properties of the electron and the probability of finding an
electron in a certain volume of the atom. This led to Schrödinger’s equation that derived
certain values for describing the likelihood of locating an electron at certain radii from the
nucleus and describing the electron’s properties in each of these volumes in 3-D space.
Erwin Schrödinger
Schrodinger’s equation for the x-axis:
𝑑2 𝛹
𝑑𝑥 2
+
8𝜋2
ℎ2
(E-V)Ψ
= 0
Solving for all three axes leads to three quantum numbers to explain the location and
properties in multi-electron atoms.
Werner Heisenberg (1901-1976)
Heisenberg stated that due to the physical methods of detection of the electron and
small particles it is impossible to be certain of both its location and momentum according
to the equation: Δx · Δ(mv) ≥ h/4π. This is Heisenberg’s Uncertainty Principle.
Current model of the atom shows the “electron cloud” as a probability distribution.
Noyau = nucleus
Probabilité de présence de l’électron = probability of the presence of the electron