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The Bohr Model
At the same time that discoveries were being made with radioactivity, physicists and chemists
were studying how light interacted with matter. These studies began the field of quantum
mechanics and helped solve the structure of the atom.
Quantum Mechanics
The current theory of atomic structure based on the wave properties of electrons; also known as
the wave-mechanical model of the atom.
Quantum Mechanics Sheds Light on the Atom: The Bohr Model
Physicists and chemists studied the nature of the light that was given off when electric currents
were passed through tubes containing gaseous elements (hydrogen, helium, neon) and when
elements were heated (e.g., sodium, potassium, calcium, etc.) in a flame. They passed the light
from these sources through a spectrometer (a device containing a narrow slit and a glass
prism).
Photo courtesy NASA
White light passing through a prism.
Photo courtesy NASA
Continuous spectrum of white light.
When you pass sunlight through a prism, you get a continuous spectrum of colors like a
rainbow. However, when light from elements such as hydrogen and helium were passed
through a prism, they found a dark background with discrete lines.
Photo courtesy NASA
Hydrogen spectrum
Photo courtesy NASA
Helium spectrum
Each element had a unique spectrum and the wavelength of each line within a spectrum had
a specific energy.
In 1913, a Danish physicist named Niels Bohr put Rutherford's findings together with the
observed spectra to come up with a new model of the atom. Bohr suggested that the electrons
orbiting an atom could only exist at certain energy levels (i.e., distances) from the nucleus, not
at continuous levels as might be expected from Rutherford's model. When atoms in the gas
tubes absorbed the energy from the electric current, the electrons became excited and jumped
from low energy levels (close to the nucleus) to high energy levels (farther out from the
nucleus). The excited electrons would fall back to their original levels and emit energy as light.
Because there were specific differences between the energy levels, only specific wavelengths of
light were seen in the spectrum (i.e., lines).
Principle of Fluorescence
1. Energy is absorbed by the atom.
The atom becomes excited.
2. The electron jumps to a higher
energy level.
3. Soon, the electron drops back to
the ground state, emitting light.
The major advantage of the Bohr model was that it worked. It explained several things:
 Atomic spectra - discussed above
 Periodic behavior of elements - elements with similar properties had similar atomic spectra.
o Each electron orbit of the same size or energy (shell) could only hold so many
electrons.
First shell = two electrons
Second shell = eight electrons
Third shell and higher = eight electrons
o When one shell was filled, electrons were found at higher levels.
o Chemical properties were based on the number of electrons in the outermost shell.
Elements with full outer shells do not react.
Other elements take or give up electrons to get a full outer shell.
Bohr's model was the predominant model until new discoveries in quantum mechanics were
made.
Electrons Can Behave as Waves: The
Quantum Model of the Atom
Although the Bohr model adequately explained how atomic spectra worked, there were several
problems that bothered physicists and chemists:
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Why should electrons be confined to only specified energy levels?
Why don't electrons give off light all of the time?
o As electrons change direction in their circular orbits (i.e., accelerate), they should
give off light.
The Bohr model could explain the spectra of atoms with one electron in the outer shell very
well, but was not very good for those with more than one electron in the outer shell.
Why could only two electrons fit in the first shell and why eight electrons in each shell after
that? What was so special about two and eight?
Obviously, the Bohr model was missing something!
In 1924, a French physicist named Louis de Broglie suggested that, like light, electrons could
act as both particles and waves. De Broglie's hypothesis was soon confirmed in experiments
that showed electron beams could be diffracted or bent as they passed through a slit much like
light could. So, the waves produced by an electron confined in its orbit about the nucleus sets
up a standing of specific wavelength, energy and frequency (i.e., Bohr's energy levels) much
like a guitar string sets up a standing wave when plucked.
Another question quickly followed de Broglie's idea. If an electron traveled as a wave, could you
locate the precise position of the electron within the wave? A German physicist, Werner
Heisenberg, answered no in what he called the uncertainty principle.
Werner Heisenberg Uncertainty Principle
It is impossible to simultaneously know exact position and speed of an electron in an atom.
Therefore, electrons should not be viewed as moving in well-defined orbits about the nucleus.
With de Broglie's hypothesis and Heisenberg's uncertainty principle in mind, an Austrian
physicist named Erwin Schrodinger derived a set of equations or wave functions in 1926 for
electrons. According to Schrodinger, electrons confined in their orbits would set up standing
waves and you could describe only the probability of where an electron could be. The
distributions of these probabilities formed regions of space about the nucleus were called
orbitals. Orbitals could be described as electron density clouds. The densest area of the
cloud is where you have the greatest probability of finding the electron and the least dense area
is where you have the lowest probability of finding the electron.
Wave Functions
The wave function of each electron can be described as a set of three quantum numbers:
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Principal number (n) - describes the energy level.
Secondary/Azimuthal number (l) - how fast the electron moves in its orbit (angular
momentum); like how fast a CD spins (rpm). This is related to the shape of the orbital.
Magnetic number (ml) - its orientation in space.
It was later suggested that no two electrons could be in the exact same state, so a fourth
quantum number (i.e. spin) was added. This number was related to the direction that the
electron spins while it is moving in its orbit (i.e., clockwise, counterclockwise). Only two
electrons could share the same orbital, one spinning clockwise and the other spinning
counterclockwise.
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Spin number (ms) – spin of the electrons in an orbital.
The orbitals had different shapes and maximum numbers at any level:
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s (sharp) - spherical (max = 1)
p (principal) - dumb-bell shaped (max = 3)
d (diffuse) - four-lobe-shaped (max = 5)
f (fundamental) - six-lobe shaped (max = 7)
The names of the orbitals came from names of atomic spectral features before quantum
mechanics was formally invented. Each orbital can hold only two electrons. Also, the orbitals
have a specific order of filling, generally: s, p, d, f….
The resulting model of the atom is called the quantum model of the atom.
Quantum model of a sodium atom.
Sodium has 11 electrons distributed in the following energy levels:
1. one s orbital - two electrons
2. one s orbital - two electrons and three p orbitals (two electrons each)
3. one s orbital - one electron
Right now, the quantum model is the most realistic vision of the overall structure of the atom. It
explains much of what we know about chemistry and physics. Here are some examples:
Chemistry
 The Periodic Table - the Table's pattern and arrangement reflects the arrangement of
electrons in the atom.
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Elements have different atomic numbers - the number of protons or electrons
increases up the table as electrons fill the shells.
Elements have different atomic masses - the number of protons plus neutrons
increases up the table.
Rows - elements of each row have the same number of energy levels (shells).
Columns - elements have the same number of electrons in the outermost energy
level or shell (one to eight).
 Chemical reactions - exchange of electrons between various atoms (giving, taking, or
sharing). Exchange involves electrons in the outermost energy level in attempts to fill the
outermost shell (i.e., most stable form of the atom).