Transcript Bohr Model of the Atom

Bohr Model of the Atom
• Why are the emission spectra of elements
not a continuous spectrum?
• In 1913, a Danish physicist named Niels
Bohr tried to discover the answer, using
hydrogen.
• Bohr proposed that hydrogen had only
certain allowable energy states.
Hydrogen’s Energy States
• The lowest energy state of an atom=
ground state.
• When the atom gains energy= excited
state.
• Bohr said hydrogen, even with only 1 e-, is
capable of many different excited states
and these are related to the orbit of the e- in
the H atom.
Hydrogen’s Energy States
• Bohr said that the single e- of hydrogen orbited
the nucleus in only certain allowed circular
orbits.
• The smaller the orbit, the lower the energy state
(energy level); the larger the orbit, the higher the
energy state (energy level)
• Bohr assigned a quantum number, n to each orbit.
The one closest to the nucleus he numbered “1,”
the next “2,” etc.
An explanation of Hydrogen’s
energy states…
• In the ground state, (n= 1) the atom does
NOT emit energy. Normal state.
• But when excited, the e- jumps up to a
higher energy level (n = 2, or 3…)
• When this excited e- drops back to the
ground state, it emits a photon equal to the
difference in energy between the 2 energy
levels.
Explanation, continued…
• Because only certain orbits are possible,
only certain frequencies of EM radiation
can be emitted.
• Think of the energy levels as rungs on a
ladder – you can only go up or down on the
rungs, NOT in-between.
ORIGIN OF THE LINES IN THE HYDROGEN
EMISSION SPECTRUM
8
7
6
5
PFUND (IR)
ENERGY
4
PASCHEN (IR)
3
2
BRACKETT (IR)
BALMER
(VISIBLE)
1
LYMAN
(ULTRAVIOLET)
The problem with Bohr…
• Bohr explained hydrogen’s spectral lines almost
perfectly.
• However, it did not explain any other element!
• It also did not explain other chemical behavior.
• And though his work was groundbreaking, later
experiments proved his model fundamentally
wrong. Electrons don’t move around the
nucleus in circular orbits.
• Alas, poor Bohr!
Enter Louis de Broglie
• In the 1920s, French graduate student
Louis deBroglie asked, “if waves (light)
can behave like particles, can particles
(electrons) behave like waves?”
The de Broglie Equation
• deBroglie eventually derived an equation to
explain the wave properties of particles.
 = _h_
mv
= wavelength of particle
h= Planck’s constant
m=mass of particle
v= velocity
The de Broglie Equation…
• This equation proves that all moving particles
have a particular wave property
• Why can’t these waves be seen, like light?
• They simply are too small. Even sensitive
instruments can’t detect these waves.
• When large objects move, when their large mass
is divided into the small number of Planck’s
constant, the resulting wavelength is very, very
tiny.
• Electrons, on the other hand, have almost NO
mass, so their wavelength is easily measured.
Heisenberg Uncertainty
Principle
• The de Broglie equation ultimately was
proven correct.
• The next scientist to contribute to our
understanding of the atom was German
physicist Werner Heisenberg.
• Heisenberg said that it’s impossible to
make a measurement on an object without
disturbing it – at least a little bit.
Heisenberg Uncertainty
Principle
• The smaller the object, the more it is disturbed by
measurement.
• For an electron, even shining light on it will
disturb its frequency and position.
• Therefore, Heisenberg concluded, “it is
impossible to know precisely both the velocity
and position of a particle at the same time.”
• Scientists of the time found it hard to accept, but
later tests proved it true.
• It shows that the uncertainty of any electron’s
position is VERY large.
Schrodinger’s Wave Equation
• In 1926, Austrian scientist Erwin Schrodinger
further refined de Broglie’s equation.
• He developed a very complex equation that
treated the hydrogen electron as a wave. It
worked perfectly for not only hydrogen, but ALL
elements.
• It is the basis for the current model for electrons
in atoms, the quantum mechanical model.
Basics of Quantum Mechanics
• The Schrodinger equation defines the basis of
energy levels and possible locations of electrons
in atoms.
• An electron’s location is simply a probability of
where it can be located at a given time.
• To make probability predictions easier, quantum
mechanics assigns numbers to energy levels and
shapes to sublevels.
Basics of Quantum Mechanics
• Principal Quantum number (n) = relative size &
energy of the atomic orbitals (where an electron
can orbit the nucleus)
• As n increases, the orbital gets larger, the espend more time away from the nucleus, & the
energy level increases.
• Therefore, the Principal Quantum number
represents Principal Energy Levels.
• The lowest energy level is assigned the number 1.
Up to 7 levels are possible for hydrogen.
Energy Levels
Sublevels
• Principal Energy levels contain sublevels (also
called orbitals).
• Energy level 1 has 1 sublevel, level 2 has 2
sublevels, 3 has 3 and so on.
• The sublevels are named for their shapes, as
follows: s=spherical; p=dumbbell; d = varying
shapes (mostly lobes); and f =varying shapes
(like flower petals, almost)
• Each orbital can hold up to TWO electrons,
maxium.