2. John Dalton Robert Andrews Millikan Niels Bohr
(1803) (1909) (1912)
Joseph John Thomas Ernest Rutherford Quantum Model
(1897) (1912) (1924)
TIMELINE
3. JOHN DALTON
- Around 1803, Dalton developed an atomic theory to explain the ratios in
which elements combine to form compounds. It was the cornerstone for
modern atomic theory.
- Five main points of Dalton’s atomic theory:
1. Elements are made of extremely small particles called atoms.
2. Atoms of a given element are identical in size, mass, and other properties:
atoms of different elements differ in size, mass, and other properties.
3. Atoms cannot be subdivided, created, or destroyed.
4. Atoms of different elements combine in simple whole-number ratios to
form chemical compounds.
5. In chemical reactions, atoms are combined, separated, or rearranged.
- He also developed an assumption based on the faith of nature’s simplicity
that when atoms combine in only one ratio, it must be presumed to be a
binary one, unless some cause appear to the contrary.
- Limitations included that nowadays, Dalton’s second and third points to the
atomic theory and proven wrong.
4. J. J. THOMSON
- In 1896 , he took the cathode ray experiments a step further by firstly
improving Perrin’s version to more clearly prove cathode rays do carry negative
charges. With this , Thomson then went on to discover the electron through his
demonstration of cathode rays responding to electrode fields just as negatively
charged particles would.
- Thomson had figured out a way to determine the charge of the mass by using
both an electric and magnetic field.
- Used mutually perpendicular electric and magnetic fields to determine the
speed of cathode rays. Then with only one field turned on, he measured the
deflection of rays. These deflections depended on magnitude of field, length of
path in the field, and the speed, mass, and charge of cathode-ray particles.
- With calculation, he found reasonably consistent values for the charge-mass
ratio, which allowed him to conclude that all cathode rays consist of identical
particles with exactly the same negative charge.
- q/m for an electron is roughly 1011 C/kg
- Reasoned electrons are much smaller than atoms.
- Put forward the daring theory that atoms were divisible and tiny particles in
cathode rays were “the substance from which all the chemical elements are built
up.” Although he was wrong about electrons being the only constituents of
atoms, recognizing electrons are subatomic particles was a major advance in
atomic physics.
- Atoms are neutral and since electrons are constituents of atoms, atoms must
contain some form of positive charge.
- Suggested that atoms might consist of electrons embedded in a blob of
massless positive charge; raisin-bun model.
- Limitations included how Thomson had made no tests to indicate that
electrons were embedded in a positive blob, but merely suggested the idea.
5. ROBERT ANDREW MILLIKAN
- In 1909, he determined the unit charge of
the electron with his oil drop experiment to
be 1.60 x 10-19.
- Millikan calculated the mass of each drop
from its diameter, then observed the motion
of the oil drops in a uniform field. By
analyzing this motion, Millikan calculated the
electric force acting on each drop. He found
that the charge of each oil drop was a
multiply of 1.60 x 10-19.
- Since others had already determined the
charge-mass ratio, Millikan could now
calculate a reasonably accurate valuable for
the mass of the electron.
- Limitations included how out of 175
measurements, he only reported 58.
6. ERNEST RUTHERFORD
- By 1909, he had shown that some radioactive elements, such as radium and
thorium, emitted positively charged helium ions, which are also known as alpha
particles and when passed through a thing sheet of mica, a beam of alpha particles
will spread out.
- He had a technique that allowed him and his assistants, Hans Geiger and Ernest
Marsden, to measure the proportion of alpha particles scattered at different angles
from various materials. They would produce a pencil-shaped beam of alpha particles
and position a thin sheet of gold foil at a right angle to the beam. Then they would
use a screen coated with zinc sulfide, which would detect the scattered particles by
letting off faint flashes of light visible with a microscope. By moving the screen and
microscope around the foil, they were able to measure the rates at which alpha
particles appear at various angles. Eventually, they concluded the positive charge in a
gold atom must be concentrated in an incredibly tiny volume, so most of gold was
actually empty space.
- When using aluminium foil instead of gold, they proved that the positive charge
and most of the mass of an atom are contained in a radius less than 10 -14.
- Discovered nucleus and disproved the raisin-bun model.
- Lead to planetary model of atom consisting of electrons orbiting the nucleus of an
atom and there being an electrostatic attraction between positive nucleus and
negative electrons, which provides the centripetal force to keep the electrons in
orbit.
- Limitations included the fact that Rutherford’s model was later adjusted by Niels
Bohr because in Rutherford’s model, the electrons should spiral into the nucleus in a
few microseconds due to a constant acceleration, which would emit electromagnetic
waves that would take energy from the orbiting electrons.
7. NIELS BOHR
- Corrected the critical flaw in Rutherford’s model.
- Focuses on the quantization of energy of electrons.
- Basic principles of Bohr’s model:
1. Electrons can orbit the nucleus only at certain specific distances
from the nucleus. These distances are particular multiples of the
radius of the smallest permitted orbit meaning the orbits in an atom
are quantized.
2. The electron’s distance from the nucleus determines both the
kinetic and electric potential energy of an electron in orbit. So forth
the energy in an electron is also quantized and each orbit
corresponds to a specific energy level for the electron.
3. Only by emitting or absorbing photons of equal energy to the
difference between energy levels can an electron move from one
energy level to another. When an electron continues to orbit at a
particular energy level, no energy is radiated. Also, since the size and
shape of the orbit remains the same and at a fixed energy level,
these orbits are often referred to as stationary states.
- Limitations included him not explaining as to why energy is
quantized, why orbiting electrons do not radiate electromagnetic
energy, why a magnetic field splits the main spectral lines into
multiple closely spaced lines, and the fact that it is not accurate for
electrons to have two or more electrons.
8. QUANTUM MODEL
- In 1924, Louis de Broglie developed his theory that particles have
wave properties. He concluded this through diffraction experiments.
So forth, the principles of interference and standing waves apply for
electrons orbiting a nucleus.
- For most sizes of orbit, successive cycles of the electron wave will be
out of phase, and destructive interference will reduce the amplitude of
the wave. For constructive interference to occur, the circumference of
the orbit must be equal to a whole number of wavelengths.
- The wave nature of matter provides a natural explanation for
quantized energy levels.
- In 1926, Erwin Schrödinger derived an equation for determining how
electron waves behave in the electric field surrounding a nucleus. The
solutions to his equation are functions that define the amplitude of the
electron wave in the space around a nucleus.
- Max Born showed that the square of the amplitude of these wave
functions at any point is proportional to the probability distribution or
orbital.
- In quantum model, electrons behave as waves, which do not have a
precise location.
-Orbitals in the quantum model show the likelihood of an electron
being at a given point, not a path they follow.
- Electrons behaving as waves rather than orbiting particles in an atom
explains why they do not radiate electromagnetic energy continuously.
9. ATOMIC SPECTRA
- The atomic spectra is a range of characteristic frequencies of
electromagnetic radiation that are readily absorbed and emitted
by an atom.
- An electron can jump from one fixed orbital to another. If the
orbital it jumps to has a higher energy, the electron must absorb
a photon of a certain frequency. If it’s a lower energy, the
electron must give off a photon of a certain frequency.
- The frequency depends on the difference in energy between
the orbitals.
- This relates to Bohr’s model because Bohr’s model describes
how in order for electrons to move from one orbital to
another, the electron must release or absorb a photon of
appropriate energy.
10. MATHEMATICS
Mass: Charge:
Thomson - Thomson -
m = Bqr/v q = Fe/E
Millikan - q = Fm/vB
m = qE/g q = vm/Br
Bohr - Millikan -
m = h2/4π2rke2 q = mg/E
m = Eh2/2π2k2e4 q = ne
Quantum Model - Rutherford -
m = nh/2πvr q = E d/kq
1 p 2
q = E d/kq
2 p 1
