Download Quantum Theory Historical Reference

Quantum Theory Historical Reference
Particle-Wave Nature of EMR and Matter
1. Isaac Newton (1644-1727):
a) classical theory of physics
b) calculus
c) optical theory – particle nature of light
2. Thomas Young:
a) diffraction pattern proved wave nature of light
- in-phase, out-of-phase waves
- constructive, destructive interference
- diffraction pattern
3. James Clerk Maxwell: EMR – energy traveling through space in the form of an
electric field perpendicular to a magnetic field. pp 180 – 183
4. Albert Einstein(1879-1955): photoelectric experiment (1905) proved particle/wave
duality of light. p 184
5. Max Plank(1858-1947): black-body experiment (1900) resulted in ultraviolet
catastrophe proving that Newtonian physics inadequate for describing nature of EMR –
spawned birth of quantum physics.
6. Louis de Broglie: Theorized that fundamental particles may also possess particlewave duality – proposed crystal diffraction pattern experiment to test theory.
7. Davisson/Germer – Thomson/Reid: Separate teams who conducted the crystal
diffraction pattern experiment resulting in proof of the particle-wave nature of
fundamental particles.
Atomic Structure
1800s
1. Sir Humphrey Davies(1778-1829): Decomposed compounds with electric current.
Theorized compounds held together by electric forces.
2. Michael Faraday(1791-1867): Determined quantitative relationship between amount
of electricity and amount of compound decomposed.
3. George Stony(1826-1911): Suggested that units of charge are associated with atoms
and proposed calling them electron.
4. JJ Thomson (1856-1940): Cathode ray experiments proved existence of electrons.
Determined charge/mass ratio of electrons: e/m = 1.75882 x 108 C/g
Plum pudding model
5. Eugen Goldstein(1850-1930): Observed positively charged particles, “canal rays,”
pass backward toward the cathode in a cathode ray tube suggesting the existence of
positively charged particles (protons) in atoms.
1900s
6. Robert Milliken (1868-1953): Oil drop experiment – charge of e-  1.60218 x 10-19C
Mass of e-  9.10940 x 10-28g equivalent to only about 1/1836 the mass of H.
Charge of one mole of electrons = 96,485 C and is equivalent to 1F (faraday)
7. Ernest Rutherford(1871-1937): Discovered alpha () and beta () particles.
Determined alpha particles to be positively charged. Gold foil experiment proved
existence of positively charged and extremely dense nucleus surrounded by clouds of
electrons at relatively large distances from the nuclei – Rutherford model. Nuclear
diameter approximately 10-5 nm or 10-14m.
8. H.G.J Moseley(1887-1915): Through the study of x-ray spectra determined that
elements differ from preceding elements (PT) by having one more positive charge on
nucleus – atomic number. Arrangement of PT according to atomic number – properties
of elements correspond to atomic number (position of table).
9. James Chadwick(1891-1974): Bombarded atoms with high speed  particles
resulting in the expulsion of neutrons from the nucleus. Thus the discovery of neutrons.
Isotopes.
10. Johann Balmer(1825-1898) & Johannes Rydberg(1854-1919): Observed emission
spectra (line spectra) of hydrogen and developed the Balmer-Rydberg equation relating
the wavelength of line spectrum to electronic transitions
1/ = R[1/(n1)2 – 1/(n2)2] R Rydberg constant: 1.097 x 107 m-1
n1 = ground state n2 = excited state
11. Neils Bohr(1885-1962): Studied H spectra (1913) and established the Bohr (H)
model of the atom, quantized energy of electrons, quantum postulates. Bohr equation
relating energy required for electron transitions
E = -B[1/(n2)2 – 1/(n1)2] B Bohr constant: 2.18 x 10-18J/eQuantum Postulates: 1. Energy of electrons quantized, electrons can reside in only
certain energy levels. 2. Electrons can transition between levels by gaining or losing
energy. 3. Energy (location) of electrons can be described by quantum numbers.
12. Louis de Broglie(1892-1987): Predicted that fundamental particles might display
wave properties under certain circumstances. His doctoral thesis (1925) proposed that a
particle of mass (m) and velocity (v) should have a wavelength associated with it.
Ultimately explains the quantized energy of electrons.
de Broglie   = h/(mv) h = Plank’s constant: 6.63 x 10-34 J.s
In order to observe the wave nature of matter, the de Broglie  must be large such that it
is measurable. Only fundamental particles (extremely small masses) have such ’s and
obey quantum mechanics. Large objects, such as a baseball, have de Broglie ’s on the
order of 10-34m and obey classical mechanics.
13 C. Davisson(1882-1958) and L.H. Germer(1896-1971): Based upon calculations of
de Broglie ’s for electrons, used crystal of Ni (empty spaces between atoms smaller than
) to diffract beam of electrons resulting in a diffraction pattern proving the wave nature
of fundamental particles. (1927)
14. Werner Heisenberg(1901-1976): Heisenberg Uncertainty Principle(1927): one
cannot determine, simultaneously, the position and momentum of a fundamental particle.
15. Erwin Schrodinger(1887-1961): Developed quantum statistical mechanics (branch
of mathematics used to decscribe the wave properties of fundamental particles) and
produced the Schrodinger equation (1926) used to determine the probable distribution of
electrons around a nucleus. Solutions to the equation provide us with the shapes and
locations of atomic orbitals. Equations applies only to 1 electron species.


 h2
( + V
82m x2
y2
z2)
16. Wofgang Pauli(1900-1958): Pauli Exclusion Principle: In order for 2 electrons to
occupy the same orbital, they must have opposite spins.