Download 4.1-Models of the Atom

• Dalton’s model proposed
that the atom was the smallest
particle of matter and was
indivisible (like billiard balls)
all matter was composed of atoms,
the atoms of each element were
different
•Thompson’s Model: The
positive charges fills the atom while
the electrons were embedded
throughout the atom (raisin bun
model)
•Thomson discovered the electron
and since the electron was
negative, but atoms neutral, there
had to positive charges in atoms
The positive fluid fills the atom while the electrons were
embedded throughout the atom (raisin bun model)
Thomson used a beam of
cathode rays in a CRT with
electric field and a magnetic
field perpendicular to the
direction of beam travel
With only the electric
field on, the beam was
deflected
• With only the magnetic field on, the cathode rays were
deflected into a curved path
 When both fields were on, and the field strengths equal,
the cathode rays were not deflected
Thomson showed that they had a negative charge
and that they had mass, so had to be particles
coming from the atom
Thomson determined the charge to mass ratio of the cathode
rays using the curvature of the beam caused by the magnetic
field
Fc  Fm
2
mv
 qvB
r
mv
 qB
r
v
q

B r m
Thomson's experiment could
not determine the mass or the
charge, only the ratio
Thompson determined the
speed of the cathode rays:
Since the beam was not
deflected when both fields
were adjusted correctly,


Fe
E
Fe  E q
q
Fm  qvB
Fe = Fm (balanced forces, 1st Law)

E q  qvB

E
v
B
Rutherford Model
•Around 1911 Rutherford, Marsden and Geiger performed
experiments to test the Thomson model
•Alpha particles from radioactive sources were directed at
thin gold foils (about 100 atoms thick)
The Thomson model predicted that most
of the alpha particles would go straight
through, and only a few would be
deflected at small angles since the
electrons in the atom have much less
mass than -particles.
Rutherford experiment
animation
•Most of the -particles went straight through undeflected,
some were deflected at angles of more than 10o and a few
were deflected almost straight back
•He concluded that most
of the atom was empty
space with most of the
mass and all of the
positive charge
concentrated in a very
small region (the nucleus)
•Scattering angles indicated the size of the nucleus was
about 10-15 to 10-14 m in radius
Rutherford's model had 2 problems:
1) The model could not
explain the light emitted
from hot gases
2) Maxwell’s theory predicted that accelerating electrons
would radiate EMR, so the electrons in Rutherford’s model
should radiate a continuous spectrum and lose energy as
they spiraled into the nucleus
•Niels Bohr developed a theory to explain
why the atom was stable:
•He suggested that the electron could not lose energy
continuously, but had to do so in “quantum jumps” and
that the electrons moved around the nucleus in circular
orbits, and only certain orbits were allowed
•Each orbit represented a certain amount of energy and that
the electron would move in the orbit without radiating energy
(although this violated classical physics)
The allowed orbits
were called stationary
states, a photon would
be absorbed when the
electron jumped from
one state to a higher
state, a photon would
be emitted if the
electron dropped to a
lower state
Bohr E-level applet
The radius of each orbit was given by the equation
r  n r1
2
radius of
Bohr orbit
number of the
orbit or energy
level: 1, 2, 3, 4,
etc.
r1 = radius of 1st orbit,
the ground state
(5.29 x 10-11 m for
hydrogen atom)
(on data sheet under “Atomic Physics”)
• Each higher orbit represents more energy.
• The energy of the electron in each orbit was found by
(on data sheet)
Energy of
electron in
energy level n
1
En  2 E1
n
Energy
level
Energy of
electron in
energy
level 1
For hydrogen, E1 = -13.6 eV or – 2.18 x 10-18 J
•The zero point for the energy is when the
e- is an infinite distance from the nucleus
•The negative sign means that the electron
needs that much energy to be removed
from the atom (to ionize the atom)
•The energy of the photon that an atom
absorbs or emits can be found by
hf = Ef – Ei
(energy in the final orbit – energy in the
initial orbit)
•When the atom is cool or unexcited, the
electron is in the ground state (n = 1)
•In a sample of excited hydrogen atoms,
each electron will release a photon as it
drops to a lower level, forming a bright line
spectrum
Where does this energy come from?
EMR - photons
• If the photon has exactly
the right amount of energy,
the atom will absorb this
energy and the photon will
no longer exist.
High speed electrons
(a la Frank Hertz exp)
• If the electron has enough
energy, the atom will
absorb (in a collision) this
energy and the remainder
would exist as the kinetic
energy of the electron
Examples
1.Determine the wavelength of the fourth line in the Balmer
series (nf = 2, ni = 6).
 = 4.09 x 10-7 m
2. If an electron absorbs exactly 2.86 eV from a photon,
what was the electron’s stationary state, and where did it
move to? When it moved back to its stationary state, what
colour was the EMR emitted?
3. Determine the energy emitted by an electron as it drops
from the 5th to the 3rd energy level of a hydrogen atom.
E = -0.967 eV
The atom lost 0.967 eV of energy.
Summary
•The Bohr model accurately explained atomic spectra (emission
and absorption)
•The energy of each orbit is quantized
•Explained why atoms are stable
•Explained the chemical and physical properties of the elements
•Could not explain why electrons could only be found in certain
orbits
•Could not explain why some lines in the emission spectra were
brighter than others
•Worked only for hydrogen and atoms with a single electron (for
example He+)