Download Chapter 15 PowerPoint

Atomic Physics – Part 1
To accompany Pearson Physics
PowerPoint presentation by R. Schultz
[email protected]
15.1 The Discovery of the Electron
Cathode rays – major item of interest near the
end of the 1800’s – what were they?
Emr? Charged particles? Neutral particles?
Path deflection caused by both electric (parallel
to field) and magnetic (perpendicular to field)
fields
15.1 The Discovery of the Electron
Conclusion: negatively charged particles
J.J. Thomson, 1897, measured charge to mass
ratio of a cathode particles using the apparatus
particles accelerated
shown:
from C to A
focused by B
deflected by magnetic
field from coil
straightened by
electric field between
D and E
15.1 The Discovery of the Electron
Thomson found q/m for cathode ray particles to
be 1.76 x 1011 C/kg
This is close to 2000 x q/m for a hydrogen ion
Thomson’s conclusion: cathode-ray particles
(today called electrons) are approximately
1/2000 the mass of a hydrogen ion (today called
protons)
15.1 The Discovery of the Electron
Since Thomson didn’t know the charge or mass
of these particles couldn’t use
qV 
1
2
mv2
to determine their speed
He balanced Fe versus Fm
was undeflected
so that particle path
15.1 The Discovery of the Electron
Then,
Fe  Fm
q E  qv B
v
E
B
Using only a magnetic field,
Fm  Fc
v2
qv B  m
r
q v

m Br
15.1 The Discovery of the Electron
Examples: Practice Problem 1, page 756
E
6.0  10 4 N C
4
v

 2.4  10
2.50 T
B
m
s
Practice Problem 1, page 758
1.0  106 m s
q
v


 1.0  108 C kg
m B r 1.0 T  0.0100 m
15.1 The Discovery of the Electron
The Mass Spectrometer uses the principles
behind Thomson’s experiment to measure the
charge to mass ratio of compounds
Read Then, Now and Future, page 759
15.1 The Discovery of the Electron
The Thomson Raisin-bun model of the atom
Recalling that Thomson concluded electrons
were approximately 1/2000 the mass of
equivalent amount of positive charge, he
concluded that atom was a mass of (+) charge,
taking up almost total volume of atom, with
tiny, near massless, electrons embedded in it
Like raisins in a bun –
Thomson Raisin Bun Model
15.1 The Discovery of the Electron
Thomson quotation
Check and Reflect page 760, questions 3, 7
SNAP booklet, page 274, questions 1, 3, 4, 7, 9,
11, 13, 16
15.2 Quantization of Charge
Millikan Oil-Drop Experiment
oil droplets sprayed into chamber
fall into region of strong electric
field
voltage across plates adjusted until
gravity is balanced
Millikan’s actual
apparatus still at CalTech
15.2 Quantization of Charge
Because of air resistance actual analysis is more
complex than that presented here
We will stick with the analysis required for this
course
15.2 Quantization of Charge
Example: Practice Problem 1, page 763
Fe  Fg
qE mg
q  5.0  105 N C  2.4  10 14 kg  9.81m s 2
q  4.7  10 19 C
4.7  10 19 C


3
e
1.60  10 19 C e
15.2 Quantization of Charge
In some problems you will need to use the
density of the oil and radius of droplet to find
its mass
4
Volume of sphere =  r 3
3
mass = density x volume
15.2 Quantization of Charge
SNAP booklet, page 280, questions 1, 2, 3, 7, 8,
10
15.3 The Discovery of the Nucleus
Discuss QuickLab 15-3, page 766
The Rutherford Experiment
15.3 The Discovery of the Nucleus
Results
most alpha particles (helium nuclei) passed
through gold foil with scattering of 1º or less
A small number were scattered backwards (at
angles greater than 140º)
15.3 The Discovery of the Nucleus
Not predicted by Thomson raisin bun model
Because of low density of mass and positive
charge in Thomson Raisin Bun model, alpha
particles (α particles) expected to blast right
through bun
Even direct collisions with electrons (raisins)
should cause little scattering
15.3 The Discovery of the Nucleus
Conclusion: Positive charge and mass not
distributed throughout whole atom, but
concentrated in a tiny fraction of total volume –
the nucleus
Atom radius approximately 10-10 m
Nucleus radius approximately 10-14 m
Atom
mostly
empty
space!
Volume to volume – like an ant in a football field
15.3 The Discovery of the Nucleus
Rutherford’s famous quotation – the scattering
“was almost as incredible as if you had fired a
15-inch shell at a piece of tissue paper and it
came back and hit you!”
15.3 The Discovery of the Nucleus
Rutherford’s Planetary Model of the atom
Note if this diagram
was to scale, nucleus
would be too small to
be visible
Review example 15.5, page 769
Do and discuss Check and Reflect, page 770,
questions 1, 2, 4, and 5
15.4 The Bohr Model of the Atom
Problem with Rutherford model:
Recall that accelerating charges produce
electromagnetic waves
Electrons orbiting nucleus undergo continuous
centripetal acceleration
Atom would give off emr, lose energy, and
electrons spiral into nucleus in small fraction of
a second!
15.4 The Bohr Model of the Atom
Fraunhofer, 1814, noticed dark lines in the solar
spectrum
Further study of the spectra led to the following
generalizations:
15.4 The Bohr Model of the Atom
• Hot dense material (e.g. a glowing solid)
produces continuous spectrum with no bright
or dark lines
• Hot gas at low pressure: emission, bright line
spectrum, characteristic of element
15.4 The Bohr Model of the Atom
• Cool gas at low pressure with light
shone through it, absorption, dark line
spectrum
Dark lines of absorption spectrum occur
at the same frequencies as the bright
line spectrum for the same gas
15.4 The Bohr Model of the Atom
Fraunhofer lines on solar spectrum later realized
to be absorption spectra of all of the gases in
the cooler outer atmosphere of the Sun
Elements identified by comparing individual
elements’ spectra with lines on the solar
spectrum
15.4 The Bohr Model of the Atom
Hydrogen’s emission spectrum had been
analyzed by Balmer, a mathematician, in 1885
Balmer noticed a pattern and formulated an
equation to predict wavelengths emitted
1
1
 1
 RH  2  2  n= 3, 4, 5, .....

n 
2
RH  Rydberg's constant for hydrogen
you won’t do
calculations
with this
formula on the
diploma exam
15.4 The Bohr Model of the Atom
Bohr realized that emitted wavelengths of light
were due to differences between quantized
energy levels in hydrogen atom
He postulated mathematically that:
• Electrons were allowed to orbit nucleus at
certain allowed radii (no others), called
stationary states, without emission of
radiation
15.4 The Bohr Model of the Atom
• Electrons in each stationary state orbit had a
fixed amount of energy – the closer to the
nucleus, the lower the energy (energy was
quantized)
• Electrons can move from one stationary state
to another by giving off or absorbing energy
in the form of EMR (remember the spectrum
of hydrogen)
15.4 The Bohr Model of the Atom
Allowed radii for hydrogen:
11
rn  5.29  10 m  n
2
Don’t need to know
for test
Allowed energy for hydrogen:

13.6 eV
En 
n2
Don’t need to know
for test
-13.6 eV is ground state (n=1) for hydrogen
n=2 and higher called “excited states”
15.4 The Bohr Model of the Atom
15.4 The Bohr Model of the Atom
View emission spectrum of hydrogen
metre stick
view hydrogen
tube through
grating
see spectral
lines on top of
metre stick
diffraction grating
hydrogen tube
15.4 The Bohr Model of the Atom
Discuss Check and Reflect, page 780, questions
2, 4, 8, 10
You should know in a non-quantitative manner,
the Bohr model of the atom: what are the basic
features, how does it get around the difficulty
in the Rutherford model, how does it explain
the spectrum of hydrogen
You should be familiar with the three types of
spectra, and how to produce each
15.5 The Quantum Model of the Atom
Bohr’s model accurately predicts hydrogen’s
spectral lines, size of the unexcited atom, and
ionization energy
Shortcomings:
Doesn’t explain why energy is quantized or how
electrons can orbit a specific radii and not emit
emr
Doesn’t work completely for anything but
hydrogen
15.5 The Quantum Model of the Atom
Doesn’t explain why spectral lines are split by a
magnetic field (Zeeman Effect)
15.5 The Quantum Model of the Atom
Recall De Broglie – previous chapter –
wavelength of particles:
h

mv
Extend the idea – the allowed orbits for
hydrogen are the right size so that each is a
whole number of “standing electron waves”
other orbital radii can’t
exist because they won’t
allow a whole # of
standing waves
15.5 The Quantum Model of the Atom
Schrödinger, 1926, wrote a wave function
(called the psi “Ψ” function) to describe
electron waves
Ψ 2 for a point is proportional to the probability
of finding an electron at that point
Orbitals – probability distributions for electrons
of different energies
15.5 The Quantum Model of the Atom
Quantum indeterminancy: in the quantum
mechanical model, electrons don’t orbit the
nucleus at defined orbits, they don’t have a
precise location
Einstein and also Schrödinger himself, had
difficulty accepting this interpretation of the
model
Interesting reading – The Schrödinger’s cat
paradox
15.5 The Quantum Model of the Atom