19.1
The Bohr Atom
p. 864
Atomic Theory before Bohr
John Dalton (1808)
Billiard-Ball Model
Atoms are the smallest particle of matter → invisible, indivisible & indestructible spheres
Atoms of each element are identicle to each other and unique from all other elements
JJ Thompson (1897) discovered “Cathode Rays” from the cathode (−) plate to the anode (+) plate
The mass of the particles (electrons) were nearly 2000 times smaller than a hydrogen atom
→ evidence that there are particles smaller than the smallest atom (H)
Plum-Pudding Model
Ernest Rutherford (1895) discovered positively charged α particles and β rays emitted by uranium
The Geiger-Marsden Experiment
• High energy α particles were directed at extremely thin gold foil surrounded by a fluorescent screen
Results → Conclusions
•
Nearly all of the α particles went nearly straight through the gold foil – less than 10° deflection → nearly all of the gold is empty space
•
The gold had a positive charge
•
The gold had no change in mass → α particles were not captured by the gold
•
Some of the α particles were deflected more than 90° → the α particles must have been repelled by a positively charged mass
Rutherford
→ negative charges were removed from the gold
developed the Nuclear Model of the atom based on the experiment of Geiger & Marsden
Nucleus = the positively charged centre of the atom having almost no volume and nearly all of the mass
Electrons = the negatively charged particles that circulate around the nucleus
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The Bohr Model of the Atom
Neils Bohr
p.870
tried to explain inconsistencies between classical mechanics and the results of :
1.
black body radiation (Kirchhoff),
2.
the photoelectric effect (Planck & Einstein),
3.
the atomic theory of Rutherford.
studied the spectrum of light emitted by hydrogen
determined the mathematical relationship that would explain the bright line spectral pattern of hydrogen
built upon the research of Balmer (visible spectrum) and Rydberg (ultraviolet)
received Nobel Prize in Physics in 1922 (age 37)
Bohr Postulates
(p.871)
1. Circular Orbits
Electrons exist in circular orbits about the nucleus. Centripetal force = electrostatic attraction between the nucleus (+) / e−
2. Energy Levels
Electrons can only orbit on certain discrete energy levels. The energy levels are quantized (limited to a set of values)
3. Electron Jumps Electron does not emit energy (light) when stays in the same orbital (energy level)
Electron can jump up to a higher energy level (excited state) if it absorbs energy equal to the difference in energy levels
Excited electron can jump down to a lower energy level by releasing photon of energy equal to the difference in energy levels.
Electron Energy Levels
Bohr related the difference in the electron energy levels (∆E) to the energy of the photons of light emitted (ħf)
| Ef – Ei | = ∆E = ħf
Bohr determined the radius of the ground state (lowest) energy level of hydrogen = Bohr radius
Bohr determined the energy levels of hydrogen using the relationship
En = −
r1 = 0.0529 nm = 5.29×10−11 m
13.6eV
n2
where n = the principal quantum number (any integer representing the nth energy level.
Spectral Series – LUV VIS PASCHEN
LUV
VIS
PASCHEN
Lyman series = Ultraviolet
Balmer series = Visible (ROYGBIV)
Paschen series = Infrared
jump down to ground state (E1)
jump down to E2
jump down to E3
|En – E1|
|En – E2|
|En – E3|
highest ∆E
small ∆E
Limitations of the Bohr Model
calculations are only valid for hydrogen spectra → failed to predict spectra for all other atoms/molecules
unable to explain bonding of atoms in molecules
no experimental evidence to support his postulates
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19.2
The Quantum Mechanical Atom
The radius and energy of the quantized electron orbitals were explained by de Broglie and Schrödinger.
de Broglie applied his idea of matter waves to the orbitting electrons. He concluded that wave bahaviour of the electrons must form a standing wave
pattern, so the orbital circumference must have an integral (whole number) of de Broglie wavelengths.
see Figure 19.11 (p878)
Quantum Numbers
n
ℓ
mℓ
ms
principal quantum number
orbital quantum number
magnetic quantum number
spin quantum number
Pauli Exclusion Principle
energy level of electron
orbital shape within the energy level)
orientation in an external magnetic field
electron acts as a spinning charge producing its own magnetic field
n = 1, 2, 3, . . .
ℓ ± s, p, d, f
mℓ = +1 or 0 or −1
ms = + ½ or − ½
two electrons in the same atom can have the same four quantum numbers
This principle helps to explain the periodic table → s-block, p-block, d-block and f-block, periods, and
magnetic properties of some elements.
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