energy and electrons rev 11/05/08 part one pisgah high school m. jones
Post on 16-Dec-2015
217 Views
Preview:
TRANSCRIPT
Energyand
Electrons
Rev 11/05/08
Part OnePart One
Pisgah High SchoolPisgah High SchoolM. JonesM. Jones
Goals for this unit – Know about:
1. the lines in the hydrogen spectrum, and Bohr’s atomic theory,
2. the arrangement of electrons in atoms, and the shape of the periodic table,
3. energy diagrams and electron configurations,
4. valence electrons and dot diagrams.
Light and Energy
Some background terms and concepts
Frequency
Symbol: f or f (or – Greek: nu)
Units: Hertz (Hz) or 1/sec
Pronounced “reciprocal seconds”
Historically, frequency was in units of “cycles per second”, but this made too much sense.
WavelengthSymbol: Greek: lambda)Units: meters (m) of radio waves in meters to
micrometers. of light in nanometers (10-9 m).
Wavelengths of light can be measured with a spectroscope.
The sine curve is often used to represent wave motion.
Look at the following graphs.
What is the relationship between frequency and
wavelength?
O n e w a v e l e n g t h
One wavelength
One wavelength
One wave-length
What is the relationship between frequency and
wavelength?
As the frequency increases, the wavelength decreases.
The relationship between frequency and wavelength can
be represented by:
λ
1f
f = frequency (lambda) = wavelength
λ
1f
Frequency is inversely proportional to wavelength
The relationship between frequency and wavelength can
be represented by:
The relationship between frequency and wavelength can
be represented by:
Frequency times wavelength equals a constant
k = f
For electromagnetic energy, the equation is:
c is the speed of light c = 3.00 x 108 m/sec
c = f
The equation can also be written as:
c = speed of light= wavelength = frequency
c =
Light is part of the electromagnetic
spectrum
Refer to your reference tables.
Electromagnetic Spectrum
Longer wavelengthLower frequencyLower energy
Shorter wavelengthHigher frequency
Higher energy
Electromagnetic Spectrum
Longer wavelengthLower frequencyLower energy
Shorter wavelengthHigher frequency
Higher energy
400 nm700 nm
The visible spectrum was discovered by …
Dr. Roy G. Biv
Red Orange Yellow Green Blue Indigo Violet
700 nmLower energy
400 nmHigher energy
Electromagnetic waves carry energy.
…inversely proportional to the wavelength.
…directly proportional to the frequency,
The energy in an electromagnetic wave is …
fhE
hc E
Energy in an EM wave
E = hc E =
hc
Small wavelength – large energy
Large wavelength – small energy
Next, a demonstration …Look at sunlight and fluorescent light
through the spectroscope.
Use the spectroscope to look at the light coming from various gas discharge tubes.
Record the colors and their order in the spectrum of hydrogen.
StopComplete the observation of
atomic spectra, then continue.
What did you see?
Hydrogen Spectrum
And now something completely different,
some history.
Thomson suggested the plum pudding model with many,
many electrons throughout
the atom.
J. J. Thomson discovered the electron in 1897.
Thomson suggested the plum pudding model with many,
many electrons throughout
the atom.
J. J. Thomson discovered the electron in 1897.
Rutherford suggested in 1911 that electrons might exist outside the nucleus in a “planetary” arrangement.
Ernest Rutherford explained that atoms had a small, dense,
positively nucleus.
Hantaro Nagaoka Japan, 1904
Enter Niels Bohr
Bohr also knew about Rutherford’s “planetary” model of the electrons.
Bohr was there right after the “gold foil” explanation was published.
Niels Bohr, with a brand new PhD in physics from U. Copenhagen, went first to Thompson, then to Rutherford to study in 1912.
Niels Bohr
In 1913 he proposed a structure of the atom with Rutherford’s nucleus and
electrons in discrete energy levels.
While at Manchester, Niels Bohr studied the spectra of elements, and how these might relate to the internal structure of atoms.
Niels Bohr
He used this new “quantum theory” to help explain how energy could be absorbed and emitted.
Niels Bohr knew of the work done by Max Planck and incorporated it into
his atomic theory.
Max PlanckIn 1900 Max Planck introduced an
unusual idea.Energy exists in “packets”, or quanta.The beginnings of Quantum Theory.A quantum of energy is the smallest
amount of energy possible.Energy exists only in multiples of
these quanta.
Cadmium Sulfide
Albert EinsteinIn 1905 Einstein used Max Planck’s
idea of quanta to explain the photoelectric effect.
Photon of light
Electron knocked loose
Gave credibility to quantum theory.
Albert Einstein
A photon of just the right energy can knock an electron out of an atom.
A photon is a “packet of light”, or quantum of energy.
Einstein explained with the Quantum Theory what could not be explained with classical physics.
Niels Bohr
… that electrons can only change energy levels when
they absorb or give off a certain amount of energy.
(1913)
Bohr said that electrons could exist only in certain discrete energy levels, and …
Hydrogen atom
nucleus
Discrete energy levels for electrons
electron
Hydrogen atom
Electrons can exist at this level,
Hydrogen atom
…or in this level,
Hydrogen atom
…or in this level,
Hydrogen atom
…but not in between the levels.
Hydrogen atom
Unless the electron is absorbing energy, or …
Giving off energy
Niels Bohr
Why ???
When observed through a diffraction grating, specific lines of color are observed.
When high voltage is connected to the hydrogen discharge tube, a
bluish light is given off.
Niels Bohr
… the electron moves up to a higher energy level.
The electron in a hydrogen atom gains energy from the electricity passing thru the tube and …
Niels Bohr
The electron in the excited state is “unstable”.
… gives off light of a certain energy and wavelength.
The electron drops to a lower energy level, and …
Don’t forget:
E = hc
= hcE
andThe wavelength is inversely proportional to the energy
Hydrogen atom
Hydrogen atom
energy
The electron absorbs energy and …
Hydrogen atom…the
electron is elevated
to the next energy level
The electron is unstable and
“wants” to return to a lower
energy level
Hydrogen atom
Hydrogen atom
Light of a particular wavelength is given off
Each line has a wavelength and color that corresponds to the difference in energy
between the two levels.
A line in the hydrogen spectrum is produced when an electron moves from higher energy level to a lower one.
Don’t forget:
E = hc
= hcE
andThe wavelength is inversely proportional to the energy
Energy and
E1
E2
E1
E2
E is large
E is small Longer
Shorter
54321
Hydrogen atom
The energy levels are
numbered …
54321
Suppose an electron is in level 5 and drops to level 2.
Hydrogen atom
The energy levels are
numbered …
54321
Then purple light with a wavelength of 434 nm will be emitted.
Hydrogen atom
The energy levels are
numbered …
54321
Hydrogen atom
Suppose an electron is in level 4 and drops to level 2.
The energy levels are
numbered …
54321
Hydrogen atom
Then you get blue-green light with a wavelength of 486 nm
The energy levels are
numbered …
54321
Hydrogen atom
Suppose an electron is in level 3 and drops to level 2.
The energy levels are
numbered …
54321
Hydrogen atom
Then red light with a wavelength of 656 nm will be emitted.
The energy levels are
numbered …
54321
The colors of the visible lines come from the energy given off by electrons moving from higher energy levels down to level 2.
Hydrogen atom
The energy levels are
numbered …
Don’t forget:
E = hc
= hcE
andThe wavelength is inversely proportional to the energy
The Hydrogen Spectrum
Each color represents the transitions of gazillions of electrons in gazillions of H atoms going from higher energy levels to the second energy level.
Why can’t we see them?
There can be transitions among seven energy levels.
There must be a lot more lines in the spectrum of
hydrogen.
Transitions in the H-spectrum7654
3
2
1
Transitions to the first energy level produce ultraviolet lines.
Transitions in the H-spectrum7654
3
2
1
Transitions to the second energy level produce visible lines.
Transitions in the H-spectrum7654
3
2
1
Transitions to the third and higher energy levels produce infrared lines.
Bohr successfully calculated the
wavelengths of all the transitions in the
hydrogen spectrum.
But only the hydrogen spectrum.
The spectra of elements with more than one
electron could not be accurately predicted.
“Regardless of it’s shortcomings and the
modifications that were later applied, Bohr’s model of the atom was the first successful attempt to make the internal structure of the atom agree with spectroscopic data.”
Asimov, 1964
The Arrangement of Electrons in the
Quantum Mechanical Model of the Atom
The Modern View of the Atom:1. A small, dense positively
charged nucleus which contains protons and neutrons.
The Modern View of the Atom:2. Electrons which exist outside of
the nucleus at …
1. various distances from the nucleus, and at …
2. various energy levels.
The Electrons3. The electrons can have both
a mass, as does matter, and a wavelength, as does light energy.
The Electrons4. The electrons themselves are
not little solid spheres in orbit around the nucleus, but exist as a “fog” of half-energy, half-matter. The electrons can behave as either matter or energy, depending on the experiment.
Energy Levels5. Based on the ideas of Bohr, the
electrons are located …a)… in major energy levels,
b)… in energy sublevels within major energy levels,
c)… in orbitals within each sublevel.
The energy levels are like an organizational chart for a
business:
e lectron e lectron
O rbita l
e lectron e lectron
O rbita l
Sublevel
e lectron e lectron
O rbita l
e lectron e lectron
O rbita l
Sublevel
M ajor energy level
Quantum Numbers1. Each electron in an atom has a set
of four (4) quantum numbers.
2. The quantum numbers are like an address: name, street, city, state.
3. Pauli exclusion principle: no two electrons in the same atom can have the same set of quantum numbers.
Quantum NumbersAddress Quntm. # Sym. What it tells
State Principal n Major energy level
City Azmuthal L Sublevel
Street Magnetic ML Orbital
Name Spin MS Which e- in orbital
Luckily, we will only deal with the Principal Quantum Number
What’s coming next?1. 2n2, and the shape of the periodic
table
2. Energy levels and sublevels
3. s, p, d, f and the periodic table
4. Orbitals, spin & energy diagrams
5. e- config., valence e-, dot diagrams
Click here to go to Part Two.
top related