Upload
selia
View
61
Download
1
Embed Size (px)
DESCRIPTION
The Quantum Mechanical Model of the Atom. Chapter 7. The nature of light. Light is electromagnetic radiation: a wave of oscillating electric & magnetic fields. Wave properties. A wave has wavelength, frequency, and amplitude. Wave properties. Wavelength = (lambda), in m - PowerPoint PPT Presentation
Citation preview
1
The Quantum The Quantum Mechanical Model Mechanical Model
of the Atomof the AtomChapter 7
2
The nature of lightThe nature of lightLight is electromagnetic radiation: a wave of
oscillating electric & magnetic fields.
3
Wave propertiesWave propertiesA wave has wavelength, frequency, and amplitude
4
Wave propertiesWave properties
Wavelength = (lambda), in mFrequency = (nu), in cycles/sec or s-1 Wavelength & frequency related by wave speed:
Speed of light c = 3.00 x 108 m s-1
c
5
Example1Example1Calculate the frequency of light with = 589 nm
6
Example 1Example 1Calculate the frequency of light with = 589 nm
= c c = 3.00 x 108 m s-1
Must convert nm to m so units agree (n = 10-9)
c
c
3.00 108 ms 1
589 10 9 m5.09 1014 s 1
7
Example 2Example 2Calculate the wavelength in nm of light with
= 9.83 x 1014 s-1
c
c
3.00 108 ms 1
9.831014 s 1 3.05 10 7 m
305 10 9 m 305nm
8
Electromagnetic spectrum = all Electromagnetic spectrum = all wavelengths of electromagnetic radiationwavelengths of electromagnetic radiation
9
Photoelectric effectPhotoelectric effectLight shining on metal surface can cause metal to
emit e- (measured as electric current)
10
Photoelectric effectPhotoelectric effectClassical theory
more e- emitted if light either brighter (amplitude) or more energetic (shorter )
e– still emitted in dim light if given enough time for e- to gather enough energy to escape
11
Photoelectric effectPhotoelectric effectObservations do not match classical predictions!
A threshold frequency exists for e- emission: no e- emitted below that regardless of brightness
Above threshold , e- emitted even with dim lightNo lag time for e- emission in high , dim light
12
Albert Einstein explainsAlbert Einstein explainsphotoelectric effectphotoelectric effect
Light comes in packets or particles called photonsAmount of energy in a photon related to its
frequency
h = 6.626 x 10-34 J s
E h hc
13
Example 3Example 3What is the energy of a photon with wavelength
242.4 nm?
E h hc
E 6.626 10 34 Js 3.00 108 ms 1
242.4 10 9 m8.20 10 19 J
14
Example 4Example 4A Cl2 molecule has bond energy = 243 kJ/mol.
Calculate the minimum photon frequency required to dissociate a Cl2 molecule.
15
Example 4Example 4243103 J
mol
1 mol6.02 1023 photons
4.04 10 19 J
1 photon
E h
Eh
4.04 10 19 J
6.626 10 34 Js6.10 1014 s 1
16
Example 4Example 4What color is this photon?
A photon with wavelength 492 nm is blue
c
3.00 108 ms 1
6.10 1014 s 1 4.92 10 7 m 492 10 9 m
17
Example 4Example 4 A Cl2 molecule has bond energy = 243 kJ/mol. Calculate
minimum photon frequency to dissociate a Cl2 molecule. What color is this photon?
A photon with wavelength 492 nm is blue
243103 Jmol
1 mol
6.02 1023 photons
4.04 10 19 J1 photon
E h
Eh
4.04 10 19 J
6.626 10 34 Js6.10 1014 s 1
c
3.00 108 ms 1
6.10 1014 s 1 4.92 10 7 m 492 10 9 m
18
Emission spectraEmission spectraWhen atom absorbs energy it may re-emit the
energy as light
19
Emission spectraEmission spectraWhite light spectrum is continuous
Atomic emission spectrum is discontinuous
Each substance has a unique line pattern
20
Hydrogen emission spectrumHydrogen emission spectrum
Visible lines at410 nm (far violet)434 nm (violet)486 nm (blue-green)656 nm (red)
21
Emission spectraEmission spectraClassical theories could not explain
Why atomic emission spectra were not continuousWhy electron doesn’t continuously emit energy as it
spirals into the nucleus
22
Bohr modelBohr modelNiels Bohr’s model to explain atomic spectra
electron = particle in circular orbit around nucleusOnly certain orbits (called stationary states) can exist
rn = orbit radius, n = positive integer, a0 = 53 pm
Electron in stationary state has constant energyRH = 2.179 x 10–18 J
Bohr model is quantized
rn n2a0
En RH
n2
23
Bohr modelBohr modele– can pass only from one allowed orbit to anotherWhen making a transition, a fixed quantum of
energy is involved
E electron E final E initial
E electron RH
n f2
RH
ni2 2.179 10 18 J 1
ni2
1n f
2
24
25
Bohr modelBohr modelCalculate the wavelength of light emitted when the
hydrogen electron jumps from n=4 to n=2
Photon energy is an absolute amount of energyElectron absorbs photon, ∆Eelectron is +Electron emits photon, ∆Eelectron is –
E electron 2.179 10 18 J 1ni
2 1n f
2
Eelectron 2.179 10 18 J142
122
4.09 10 19 J
Ephoton 4.09 10 19 J
26
Bohr modelBohr modelCalculate the wavelength of light emitted when the
hydrogen electron jumps from n=4 to n=2
486 nm corresponds to the blue-green line
E electron 2.179 10 18 J 1ni
2 1n f
2
Ephoton 4.09 10 19 J
Ephoton 4.09 10 19 J hc
6.626 10 34 Js 3.00 108 ms 1
4.86 10 7 m 486 10 9 m 486nm
27
ExampleExampleWhat wavelength of light will cause the H electron
to jump from n=1 to n=3? To what region of the electromagnetic spectrum does this photon belong?
E electron 2.179 10 18 J 1ni
2 1n f
2
28
ExampleExampleWhat wavelength of light will cause the H electron
to pass from n=1 to n=3?
Eelectron 2.179 10 18 J112
132
1.94 10 18 J
Ephoton 1.94 10 18 J
E electron 2.179 10 18 J 1ni
2 1n f
2
29
ExampleExampleWhat wavelength of light will cause the H electron
to pass from n=1 to n=3?
The atom must absorb an ultraviolet photon with = 103 nm
Ephoton 1.94 10 18 J hc
6.626 10 34 Js 3.00 108 ms 1
1.0310 7 m 10310 9 m 103nm
E electron 2.179 10 18 J 1ni
2 1n f
2
30
Light has both particle & wave behaviorsLight has both particle & wave behaviors
Wave nature shown by diffraction
31
Light has both particle & wave behaviorsLight has both particle & wave behaviors
Particle nature shown by photoelectric effect
32
Electrons also have wave propertiesElectrons also have wave propertiesIndividual electrons exhibit diffraction, like waves
How can e– be both particle & wave?
33
ComplementarityComplementarityWithout laser, single e–
produces diffraction pattern (wave-like)
With laser, single e– makes a flash behind one slit or the other, indicating which slit it went through –– and diffraction pattern is gone (particle-like)
We can never simultaneously see the interference pattern and know which slit the e– goes through
34
ComplementarityComplementarityComplementary properties exclude each other
If you know which slit the e– passes through (particle), you lose the diffraction pattern (wave)
If you see interference (wave), you lose information about which slit the e– passes through (particle)
Heisenberg uncertainty principle sets limit on what we can know
² x m² v h4
35
IndeterminancyIndeterminancy
Classical outcome is predictable from starting conditions Quantum-mechanical outcome not predictable but we can
describe probability region
36
Electrons & probabilityElectrons & probabilitySchrodinger applied wave mechanics to electrons
Equation (wave function, ) describe e– energyEquation requires 3 integers (quantum numbers)Plot of 2 gives a probability distribution map of e–
location = orbitalSchrodinger wave functions successfully predict
energies and spectra for all atoms
37
Quantum numbersQuantum numbersPrincipal quantum number, n
Determines size and overall energy of orbitalPositive integer 1, 2, 3 . . .Corresponds to Bohr energy levels
38
Quantum numbersQuantum numbersAngular momentum quantum number, l
Determines shape of orbitalPositive integer 0, 1, 2 . . . (n–1)Corresponds to sublevels
l
letter
0 s1 p2 d3 f
39
Quantum numbersQuantum numbersMagnetic quantum number, ml
Determines number of orbitals in a sublevel and orientation of each orbital in xyz space
integers –l . . . 0 . . . +l
40
41
Shapes of orbitalsShapes of orbitalss orbital (l = 0, ml = 0)
p orbitals (l = 1, ml = –1, 0, +1)
42
Shapes of orbitalsShapes of orbitalsd orbitals (l = 2, ml = –2, –1, 0, +1, +2)
43
What type of orbital is designated by each set of quantum numbers?n = 5, l = 1, ml = 0n = 4, l = 2, ml = –2n = 2, l = 0, ml = 0
Write a set of quantum numbers for each orbital4s3d5p
44
What type of orbital is designated by each set of quantum numbers?n = 5, l = 1, ml = 0 5pn = 4, l = 2, ml = –2 4dn = 2, l = 0, ml = 0 2s
Write a set of quantum numbers for each orbital4s n = 4, l = 0, ml = 03d n = 3, l = 2, ml = –2, –1, 0, +1, or +25p n = 5, l = 1, ml = –1, 0, or +1
45
46
Electron configurationsElectron configurations
Aufbau principle: e– takeslowest available energy
Hund’s rule: if there are 2 or more orbitals of equal energy (degenerate orbitals), e– will occupy all orbitals singly before pairing
Pauli principle: Adds a 4th quantum number, ms (spin) No two e– in an atom can have the same set of 4 quantum
numbers ⇒ 2 e– per orbital
ms 12 or 1
2