UNIT 1: STRUCTURE AND PROPERTIES
ATOMIC THEORY
Development of the Bohr Model of the Atom
Early Atomic Theory
• Ancient Greeks (e.g. Democritis) proposed that when
matter is divided into smaller and smaller pieces, a finite
limit known as the atom is ultimately reached.
• However the theory that matter was composed of 4
classical elements was accepted until the late Middle
Ages.
Dalton’s Atomic Theory (1803)
Dalton revived the atomic theory to explain several
chemical laws:
• Law of Conservation of Mass (Lavoisier, late 1700’s)
• Law of Definite Proportions (Proust, 1799)
• Law of Multiple Proportions (Dalton, 1803)
This led to the “Billiard Ball” model of indivisible
atoms of elements. These combine to form all
known compounds.
Thomson’s Model of the Atom (1897)
• Michael Faraday and Svante Arrhenius studied
electricity and batteries, suggesting that electric
charges are part of matter.
• Thomson detected and measured the mass of a
beam of negative particles in a vacuum tube (or
cathode ray tube, CRT) .
• He called the tiny particles that emerged from a
metal cathode electrons.
Thomson’s Plum Pudding Model
• Thomson explained electrical conduction in metals by the
movement of subatomic electrons in the solid.
• Electrical conduction in solutions (e,g, batteries) was
explained by the existence of charged atoms called ions.
Rutherford’s Atomic Model
• The discovery of radioactivity (,  and  rays) allowed
Ernest Rutherford to probe inside the atom.
• Based on Thomson’s model, Rutherford proposed that
alpha rays (high energy helium ions, He2+) should pass
through the positive pudding in a very thin sheet of gold
foil.
• The results were slightly different…
Gold Foil Experiment
Expected result:
Actual result:
Rutherford’s Planetary Model (1911)
• Rutherford concluded that the atom consists of a
very tiny, dense nucleus composed of the positive
charge.
• >99.99% of the atom consists of empty space.
• Tiny electrons are found orbiting the nucleus.
• A teaspoon on the atomic nuclei would weigh six
billion tonnes (6 x 1012 kg)!!!!
900 x
• In 1914, Rutherford proposed that there existed a
positively charged particle called a proton.
• In 1932, Chadwick & Rutherford found that the
nucleus also contained a neutral particle called a
neutron.
Summary of the Subatomic Particles
Particle
Location
Charge
Mass (amu)
Proton
Nucleus
+1
1.0073
Neutron
Nucleus
0
1.0087
Electron
Outside Nucleus
-1
0.00055
amu = atomic mass unit = 1.66 x 10-24 g
Problems with the Planetary Model
1) According to classical physics, electrons
orbiting the nucleus should lose energy and
emit light. This loss of energy would cause the
electrons to spiral into the nucleus, resulting in
the collapse of the atom.
2) Excited electrons should emit a continuous
spectrum of white light when they are excited.
Instead, the emission spectrum of elements are
all unique.
Emission and Absorption Spectra
• Neils Bohr observed the line spectra produced by the
excitation of gas state elements.
• Excited gases produce a unique emission spectrum.
• Cold gases will absorb produce an absorption spectrum.
The Wave Theory of Light
Light is a form of electromagnetic radiation.
All electromagnetic radiation is made up of electric
and magnetic fields. These fields oscillate in a
wave pattern as electromagnetic radiation moves
through space.
Electromagnetic radiation travels at a constant
speed in a vacuum, but the wavelength and
frequency varies.
Basic wave terms:
c= speed of light = 3.00 x 108 m/s
= lambda = wavelength (m)
f = frequency (cycles per second = s-1 = hertz)
c

f
short wavelength (blue)
long wavelength (red)
direction of movement
The Spectrum of Electromagnetic Radiation
Note: 1 nanometer (nm) = 1 x 10-9 m
The “visible” part of the spectrum is between 390 nm and 750 nm.
Name of Radiation
Radio Waves
Microwaves
Radar Waves
Infrared Light
Visible Light *
UV Light
X-Rays
Gamma Rays
Colour
Red
Orange
Yellow
Green
Blue
Indigo
Violet
Wavelength (m)
Frequency (c/s or Hz)
102
106
10-7
Increasing
Wavelength
10-14
1014
Increasing
Frequency
1022
Wavelength (m)
6.5 x 10-7
Frequency (c/s or Hz)
4.6 x 1014
4.1 x 10-7
7.3 x 1014
Note that the wavelength and frequency or inversely proportional:
If  , then f 
c

f
c  f
Planck’s Quantum Hypothesis (1900)
• Max Planck realized that the spectrum produced
by white light produced was not continuous.
• Planck proposed that light was composed of
small packets or quanta of energy called
photons.
• Different colours of light are composed photons
with different “quantums” of energy.
• The energy of a particular photon is proportional
to the frequency of the radiation.
The Particle Theory of Light
The energy of a photon is some multiple of an energy quantum called
Planck’s constant.
E  hf
Where:
E = energy of a single photon (kJ)
f = frequency (cycles/s, Hz, s-1)
h = Planck’s Constant = 6.63 x 10-37 kJs
Note: The energy of a mole of photons can be found using:
E  hfN A
Sample Wave and Particle Theory Problems:
1)
What is the wavelength of FM 92.5? (92.5 FM has a
frequency of 99.1 MHz).
2)
What is the energy of 1 photon of electromagnetic
radiation emitted by FM 92.5??
3)
What is the energy of 1 mole of these photons?
Answers: 3.24 m; 6.13 x 10-29 kJ/photon; 3.69 x 10-5 kJ/mol
Bohr’s Model of the Atom (1913)
Bohr developed a model of the atom which explained the
line spectrum of hydrogen and why the atom doesn't
collapse. His theory is made up of two postulates:
Postulate 1:
Electrons can only move in certain fixed orbits. Each orbit
corresponds to a specific energy level and an electron can
move within an orbit without losing any energy.
Postulate 2:
An electron can only move from one orbit (or energy level)
to another when it gains or loses energy.
The Hydrogen Spectrum Explained
In a discharge tube, the electrons become temporarily
excited and thus move to orbits that are farther from the
nucleus. These transitions are only temporary, though, so
the electrons return to their lowest (or ground) states.
When they do so, they give off energy.
Why Only Four Visible Lines?
Since electrons can only undergo transitions between
certain specific energy levels, only certain quantities
(quantums) of energy are given off. Since the colour of light
corresponds to the energy of possessed by a quanta (or
photon), only certain coloured lines are observed.
Animation of the Bohr Model of the Atom (Emission)
Absorption in the Bohr Model
Another Animation
• Bohr concluded that the permitted energy levels (n)
correspond to quantized transitions; only certain energy
changes are permitted as determined by the Rydberg
formula.
CALCULATING TRANSITION ENERGY:
The relative energy of any orbit (n) of the hydrogen atom can be
calculated by:
 1312
En 
n2
where En is the energy of any orbit (kJ/mol).
The transition energy (energy gained or lost as electrons change
energy levels) can therefore be calculated. This is equal to the energy
emitted by 1 mole of excited hydrogen atoms.
Etransition  En f  Eni
 Etransition
  1312    1312 



2
 n
  n2 
f
i


 
Significance of the Bohr Model
• The Bohr model explained why different elements have
unique spectra as the value of En depends on the charge
of the nucleus.
• Bohr’s model also explains Mendeleev’s Periodic Law as
the chemical and behaviour of elements is related to the
filling of Bohr’s energy levels (2, 8, 8,18 etc).
Problems with Bohr’s Model
• Calculations work for 1 electron systems (H, He+, Li2+).
• Calculations for other atoms do not agree with the
experimental results.
• The visible spectral lines of hydrogen can be split by
electric or magnetic fields fields.
• In conclusion, the Bohr model was a great leap in the
understanding of the atom. It was the first step in
quantum mechanics, the current atomic theory.
Sample Problem:
Calculate the frequency and wavelength of electromagnetic
radiation emitted from a hydrogen atom if an electron
moves from the third energy level to the first energy level.