Download The Bohr Model of the Atom

The Bohr Model of the Atom
Atomic Model Revisions
• Dalton’s model of the atom assumed the atom
to be an indestructible mass. This explained
the nature of chemical reactions to a large
degree.
• But the idea of indivisible atoms was
shattered by Thomson and Rutherford.
• Thomson came up with his own theory, called
the plum pudding model, which assumed that
electrons were imbedded in a positively
charged matrix.
• This model addressed some, but not all, of the
electrical properties of atoms.
Further Revisions:
The Rutherford Model
• Rutherford’s model of the atom composed of a
positively charged nucleus surrounded by
electrons explained a few properties of atoms,
but not all.
• In particular, it did not explain why many atoms
emit light of specific frequencies when heated.
• Rutherford also could not explain why electrons
did not collapse into the nucleus.
Big Problem
• Neither Rutherford nor Thomson’s models
could explain why the electrons of elements
and compounds emit specific wavelengths of
light, rather than a complete spectrum.
• This “bar code” is unique for hydrogen. The
lines pictured above are known as the “Balmer
series.”
• Hydrogen has several other series of lines that
are outside the visible range.
• For most of the 19th century, spectroscopists
puzzled over these lines for a wide variety of
substances, with little to show for it.
Emission Spectra of Elements
• Each element gives
off a unique
selection of
wavelengths.
• Only the visible
spectrum is shown
here! There are
more lines than
these.
Lab 4 Emission Spectra of Elements
• Look at Ne and Ar spectrum tubes with your
spectroscope. Draw one set of emission
spectra in the boxes below.
Neon
Wavelength
Argon
Wavelength
Argon and Neon Spectra
• On the diagrams on the following page, draw
a thick line with an appropriate colored pencil
through each of the major peaks that
correspond to the lines you see with the
spectroscope.
• Estimate the wavelengths, in nanometers, that
correspond to at least some of the lines that
you see using your spectrograph.
Neon Spectrum
Argon Spectrum
Questions
• How do your drawings of the emission spectra
compare to the emission spectra on the
following page?
• Do you have the same lines, or are some
missing?
• Is the spacing between the lines similar, or
different?
Identifying Elements Using Spectra
• Using the data on the back, how could
someone distinguish between samples of Ne
and Ar gases without observing them directly
in a cathode ray tube?
Atomic Theory Test
Multiple Choice Breakdown
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Rutherford model:
Nuclear structure:
What is the atomic number?
What is the mass number?
How many neutrons?
How many electrons?
Atomic mass definition
Atomic mass calculation
4 questions
4 questions
8 questions
2 questions
3 questions
2 questions
3 questions
3 questions
Modern atomic theory begins
with Niels Bohr (1885-1962).
In 1913, Niels Bohr explained why the lines in the
hydrogen spectrum are arranged the way they are.
His explanation places electrons in stable orbits
which he called "stationary states." These states
violated well-known laws of science.
Bohr saw spectrum lines as the energy difference
resulting from the movement of an electron from a
higher-energy stationary state to a lower-energy
one.
• Bohr proposed that electrons are arranged in
specific circular paths around the nucleus – a
planetary model.
• His major contribution was to propose that
electrons in a particular orbit have a fixed
energy, thus avoiding falling into the nucleus.
• The energy level of
an electron is the
region around the
nucleus where the
electron is likely to
be moving.
• According to Bohr,
electrons cannot
exist between the
energy levels – they
have to “jump” from
one level to another.
• To do this, they must
gain or lose a specific
amount of energy.
• A quantum of energy
is the amount of
energy required to
move an electron
from its present
energy level to the
next higher one.
• The energies of
electrons are
quantized. The term
“quantum leap”
comes from this idea.
• The quanta gained
or lost by every
electron can differ.
• Energy levels in an
atom are not
equally spaced.
• As the distance
from the nucleus
increases, the
energy levels are
more closely
spaced.
• The spectrum of hydrogen
contains three major sets of
spectral lines:
• The Lyman series
(ultraviolet)
• The Balmer series (visible)
• The Paschen series
(infrared)
• When the electron is in its lowest energy state,
n=1, the “ground state.”
• The electron absorbs a quanta of energy, raising it
to an excited state where n=2,3,..6.
• You can’t “see” absorption, only emission
• The absorption spectra have missing lines
More on the Bohr Model
• The same quanta are emitted when the
electron drops from the excited state to the
ground state.
• Only electrons that “fall” from higher to lower
energy states emit light.
• Bohr accounted for the hydrogen spectrum by
recognizing the lines represented those drops
into lower energy states.
Hydrogen Spectrum
• The Lyman series (ultraviolet) n>1 drops to n=1
• The Balmer series (visible) n>2 drops to n=2
• The Paschen series (infrared) n> 3 drops to n=3
• Note that other series for n>3 also exist. There is
a limit to each series because really excited
electrons escape the hold of the nucleus.
• Bohr’s theory did not explain the spectra of any
other element, or any compound.
Bohr Model Limitations
• Bohr’s notion of energy levels of electrons
resulted in a complete explanation for a single
element – hydrogen.
• Why was he able to explain only hydrogen? It
only has one electron. He could also explain a He+
ion or Li2+ ion, which also have only one electron.
• When multiple electrons are involved, they
interact with each other in complicated ways –
called the “three-body problem.”