Atomic Structure I It’s not about Dalton anymore… .
-
Upload
veronica-rawdon -
Category
Documents
-
view
219 -
download
2
Transcript of Atomic Structure I It’s not about Dalton anymore… .
Atomic Structure IIt’s not about Dalton
anymore…
http://plus.maths.org/latestnews/may-aug07/strings/atoms.jpg
First…
• To understand the electronic structure of the atom we need to review the properties of electromagnetic radiation.
Figure 7.1
Frequency and
Wavelength
c = wavelength
frequency
C = speed of light
The Wave The Wave Nature Nature of Lightof Light
Amplitude (intensity) of a wave.
The waveheight or amplitude determines radiation intensity. The wavelength is related to the energy of the radiation.
λ, ν, and Energy
• As λ decreases and ν increases, what happened to the energy of the radiation?
where h = Planck’s constant
(6.626 × 10-34 m2 kg/s)
h cE = h =
Regions of the electromagnetic spectrum.
The infinite number of wavelengths of electromagnetic radiation have been classified into groups as shown below.
SOLUTION:
Interconverting Wavelength and Frequency
Use c =
10-2
m1 cm
10-9
m1 nm
= 1.00x10-10 m
= 325x10-2 m
= 473x10-9
m
=3x108
m/s1.00x10-10
m
= 3x1018 s-
1
=
=
3x108
m/s325x10-2 m= 9.23x107 s-
1
3x108
m/s473x10-9 m= 6.34x1014
s-1
PROBLEM: A dental hygienist uses x-rays (= 1.00A) to take a series of dental radiographs while the patient listens to a radio station ( = 325 cm) and looks out the window at the blue sky (= 473 nm). What is the frequency (in s-1) of the electromagnetic radiation from each source? (Assume that the radiation travels at the speed of light, 3.00x108 m/s.)
o
325 cm
473nm
1.00Ao 10-10 m
1Ao
SOLUTION:
Calculating the Energy of Radiation from Its Wavelength
PROBLEM: A cook uses a microwave oven to heat a meal. The wavelength of the radiation is 1.20cm. What is the energy of one photon of this microwave radiation?
After converting cm to m, we can use the energy equation, E = h combined with = c/ to find the energy.
E = hc/
E =6.626X10-34J*s 3x108m/
s1.20cm
10-
2mcm
x= 1.66x10-23J
Light is a wave…right?
• Light falling on alkali metals causes electrons to be released from the metal.
• The # of electrons depends on the intensity of light.
• There are specific wavelengths of light that cause the release of e-.
• This is called the photoelectric effect.
Light is a wave…right?
• Einstein’s interpretation of the photoelectric effect (1905) was that light is quantized in packets of set energy called photons. (He won the Nobel Prize for this.)
• This meant that light had characteristics of particles!
Electrons are particles…right?
• In 1925, de Broglie stated that all particles have a wavelength described by the equation:
λ = h/p where p= momentum• Electrons show diffraction pattern
when passing through a slit• So light and particles have a dual nature.
Back to atomic structure…
• We already know an atom contains a nucleus with p+ and no. Electrons orbit the nucleus.
• It was known that atoms emit a unique spectrum of lines when excited. Rydberg derived an equation that related the lines.
• R is the Rydberg constant = 1.096776x107 m-1
= RRydberg equation -1
1
n22
1
n12
http://hyperphysics.phy-astr.gsu.edu/hbase/quantum/atspect2.html
Atomic emission spectra
Clockwise from lower left:neon, helium, hydrogen,mercury, nitrogen
Spectra Site
• http://jersey.uoregon.edu/vlab/elements/Elements.html
Absorption and emission spectra for element arranged on the periodic table
Back to atomic structure…
• Bohr theorized that the emission spectra of atoms described by Rydberg’s equation were caused by the transition of electrons between specific energy levels (orbits).
• http://www.upscale.utoronto.ca/GeneralInterest/Harrison/BohrModel/Flash/BohrModel.html
Electron locations
• When an electron occupies its usual energy level it is in the ground state.
• When an electron absorbs a photon and moves to a higher energy level it is in an excited state.
• The energy levels are “quantized”. Atoms can only transition between set levels.
• Why are the levels set where they are?
More on electrons as waves
• Since electrons have wave motion Schrödinger applied the classic wave equations to the motion of a hydrogen electron. Certain wavelengths reinforced each other and were allowed.
• This generated regions occupied by an electron of set energy termed orbitals.
More on electrons as waves
• Heisenberg stated that in measuring the electron there is uncertainty so we can only calculate a probable location for the electron. This is called the Heisenberg Uncertainty Principle.
CLASSICAL CLASSICAL THEORYTHEORYMatter
particulate,
massive
Energy continuou
s, wavelike
Since matter is discontinuous and particulate perhaps energy is discontinuous
and particulate.Observation
Theory
Planck: Energy is quantized; only certain values allowed
blackbody radiation
Einstein: Light has particulate behavior (photons)
photoelectric effect
Bohr: Energy of atoms is quantized; photon emitted when electron changes orbit.
atomic line spectra
Summary of the major observations and theories leading from classical theory to quantum theory.
Since energy is wavelike perhaps matter is wavelike
Observation
Theory
deBroglie: All matter travels in waves; energy of atom is quantized due to wave motion of electrons
Davisson/Germer: electron diffraction by metal crystalSince matter has mass perhaps energy has
massObservation
Theory
Einstein/deBroglie: Mass and energy are equivalent; particles have wavelength and photons have momentum.
Compton: photon wavelength increases (momentum decreases) after colliding with electron
QUANTUM THEORY
Energy same as Matterparticulate, massive, wavelike