Download Intro to Bohr Model - Blogs at UMass Amherst

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Introduction to the Bohr Model of Atomic Structure.
In the early 1900’s, thanks in part to Rutherford’s famous gold foil experiment, scientists knew several facts about
the structure of atoms.
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Most of the mass of an atom occupies a tiny but extremely dense and positively charged region in the
center of the atom called the nucleus. This discovery led to the recognition of protons as nuclear material;
neutrons were discovered some years later…
The much less massive electrons are found outside the nucleus, in a relatively much larger region of what is
mostly empty space.
This image is not to scale. If the nucleus were really as big as you see
it here, then the distance between the nucleus and the electrons would be 100 meters or more further away! Fact:
The total volume of an atom is mostly empty space.
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Rutherford’s nuclear atomic model was consistent with all of the available evidence known about atoms
and chemical reactions at that point in history.
A major flaw in this model, however, was that it did not explain the periodic properties of the elements first
identified by Russian scientist Dmitri Mendeleev (remember he’s the guy who created the first periodic
table). It also seemed to defy known physical laws, which predicted that electrons should be drawn into the
nucleus. Clearly a better model of atomic structure was needed.
In 1913, a scientist named Neil’s Bohr proposed an improved model of atomic structure, which was successful
in describing the line spectrum of hydrogen, shown below.
The four colors represent different wavelengths of electromagnetic radiation, which are emitted by hydrogen
gas when it is placed in a strong electric field. Bohr interpreted this observation by concluding that hydrogen’s
single electron could exist only at a limited number of distances from the nucleus. Bohr assumed that the
electron revolved around the nucleus in a circular orbit. When the electron absorbed a certain quantity of
energy, it would jump to a higher allowed orbit. These different orbits are also called energy levels. When an
electron is found in the smallest allowed orbit (closest to the nucleus), we say the electron is in the ground state.
When an electron is found in a higher energy orbit, we say the electron is in an excited state. The colors we see
in hydrogen’s line spectrum occur when an electron relaxes from a higher energy orbit to a lower energy orbit,
emitting energy as photons of light.
Today chemists and physicists recognize that Bohr’s assumption of circular orbits is incorrect. The motion of
electrons in atoms is much more complex than simple circular orbits. Indeed electrons, like light, have a dual
nature exhibiting properties of both particles and of standing waves. For this reason, energy level is a better
word choice than orbit. However we often still adopt this circular orbit assumption because its simplicity makes
the Bohr model of the atom easy to work with.
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Introduction to the Bohr Model of Atomic Structure.
The following diagrams show the wavelengths of light (EMR) produced when the excited electron in a
hydrogen atom relaxes to a lower energy level. Notice the some relaxations produce UV light, and infrared
radiation, in addition to the four visible wavelengths.
Bohr Model (circular orbit) diagram
Energy Level Diagram-notice the spacing (∆E) between
levels gets smaller as the energy level # increases. This
important point is not made clear in the Bohr diagram on
the left.
Bohr’s conclusions about hydrogen’s line spectrum were applied to the atoms of other elements, but with less
success. Soon an even better model of atomic structure was developed to replace the Bohr model. The most
modern model of atomic structure is called the Quantum Mechanical Model, which builds off of Bohr’s ideas
but is far more complex and mathematical (and beyond the scope of most first year high school chemistry
classes).
Still, his concepts do help to understand some properties of the elements, especially those with atomic numbers
1-18. One additional conclusion that is still accepted today has to do with the number of electrons a given
energy level can hold. As the energy levels increase in size (and energy, as well as distance from the nucleus),
the number of electrons an energy level can hold increases according to the formula 2(N) 2, where N is the
number of the energy level.
For example, the 1st lowest energy level (N=1) can hold 2(1) 2 = 2 electrons.
The second energy level (N=2) can hold 2(2 2) = 8 electrons.
How many electrons can the 3rd energy level hold?
When sketching Bohr diagrams of atoms with all of their electrons in the ground state, electrons always fill
lower energy levels before occupying larger, more energetic energy levels.
Here is an example, the Bohr atomic diagram for Sodium:
See if you can sketch a Bohr model of
the most common isotope of Chlorine
here: