Transcript ch 6 ppt - mvhs

Electronic Structure of Atoms
(i.e., Quantum Mechanics)
Brown, LeMay Ch 6
AP Chemistry
Monta Vista High School
1
What does light have to do with
the atomic model?
Scientists knew the nature of light but
knew little about the nature of
matter. To understand the nature of
matter, scientists studied the changes
caused in light by interaction of
matter. From these studies, scientist
tried to extrapolate information
about the nature of matter.
2
6.1: Light is a Wave

Electromagnetic spectrum:
A form of radiant energy (can travel without matter)
Both electrical and magnetic (properties are
perpendicular to each other)
http://imagine.gsfc.nasa.gov/Videos/general/spectrum.mov

Speed of Light: c = 3.0 x 108 m/s (in a
vacuum)
http://www.astronomynotes.com/light/s3.htm
Wavelength (l): distance between wave peaks
(determines “color” of light), measured in nm, m
etc.
Frequency (n): # cycles/sec (measured in HzHertz, hz= cycles/s cor= 1/s)
ln
3
6.2: Light is a Particle (Quantum Theory)

Blackbody radiation:
* Blackbody: object that absorbs all EM radiation that strikes it; it can radiate all
possible wavelengths of EM; below 700 K, very little visible EM is produced;
above 700 K visible E is produced starting at red, orange, yellow, and white
before ending up at blue as the temperature increases
◦ discovery that light intensity (energy emitted per unit of time) is proportional to
T4; hotter = shorter wavelengths
“Red hot” < “white hot” < “blue hot” Interactive Link
• Planck’s Theory: (explained blackbody radiation by quantization of energy
transfer)
Blackbody radiation can be explained if energy can be released or
absorbed in packets of a standard size he called quanta (singular:
quantum).
E  hn 
hc
l
where Planck’s constant (h) = 6.63 x 10-34 J-s
E  hn 
Animation Link
hc
Max Planck
(1858-1947)
l
4
The Photoelectric Effect

Spontaneous emission of e- from metal
struck by light; first explained by Einstein
in 1905
A quantum strikes a metal atom and the energy is
absorbed by an e-. If the energy is sufficient, e- will leave
its orbital, causing a current to flow throughout the
metal.
To explain photoelectric effect, quantization of light was
put forth by Einstein. Animation
Albert Einstein
(1879-1955)
5
6.3: Bohr’s Model of the H Atom (and only H!)


Applied quantization of energy transfer to the atomic model
Studied atomic spectrum of H to come up with atomic
model.
Atomic emission spectra:


Most sources produce light that contains many wavelengths at
once. Animation
However, light emitted from pure substances may contain only
a few specific wavelengths of light called a line spectrum (as
opposed to a continuous spectrum). Animation
Atomic emission spectra are inverses of atomic absorption
spectra.
6
Atomic Emission Spectra of C and H
Hydrogen: contains 1 red, 1 green, 1 blue and 1 violet.
Carbon: Contains many more emission lines as compared to H. Why?
7
Niels Bohr theorized that e-:
◦ Travel in certain “orbits” around the nucleus, or, are only
stable at certain distances from the nucleus
◦ If not, e- should emit energy, slow down, and crash into the
nucleus.
Allowed orbital energies are defined by:
 RH
 2.178 10
En 

2
n
n2
18
principal quantum number (n) = 1, 2, 3, 4, …
Rydberg’s constant (RH) = 2.178 x 10-18 J
Johannes Rydberg
(1854-1919)
Niels Bohr
(1888-1962)
8
Think, Pair, Share Activity
With your elbow partner, describe
Electromagnetic radiation, blackbody
radiation, Plank’s theory and
Photoelectric effect. Address each of the
above in the following terms:
1. What is it?
2. Why was it important?
3. What existing theory or concept, it
approved/disapproved.
•
9
5
4
E3
3
E2
2
E1
1
Principal Quantum Number, n
Increasing Energy, E
E5
E4
As n approaches ∞, the e- is essentially removed from the atom, and
E∞ = 0.
•
•
ground state: lowest energy level in which an e- is stable
excited state: any energy level higher than an e-’s ground state
 1
1 
E   R H  2  2 
ni 
 nf
E  R H  1
1 
 2 2
n


h
h  n f
ni 
Phased out!!
E R H  1
1 
 2 2
n


h
h  n i
nf 
ni = initial orbital of enf = final orbital of e- in its transition
Movie on e transition
11
5
4
3
2
Friedrich
Paschen
(1865 - 1947)
n
Theodore
Lyman
(1874 - 1954)
Johann
Balmer
(1825 – 1898)
Phased out!
1
Frederick
Brackett
(1896 – 1988)
?
Figure 1: Line series are transitions from one level to
another.
Series
Transition down to (emitted)
or up from (absorbed)…
Type of EMR
Lyman
1
UV
Balmer
Paschen
Brackett
2
3
4
Visible
IR
Far IR
6.4: Matter is a Wave
Planck said:
E=hc/l
Einstein said:
E = m c2
Louis DeBroglie said (1924):
h c / l  m c2
h/lmc
Louis
de Broglie
Therefore:
(1892 - 1987)
Particles (with mass) have an
m = h / cl
associated wavelength
Waves (with a wavelength) have an
l  h / mc
associated mass and velocity
13
Neils Bohr Model: Partner Activity

On a sheet of paper, take turns with your partner
drawing Bohr’s model of atom.
Draw the following in context of Bohr’s Model:
1.nucleus
2.energy levels (1,2,3,4)
3.an electron in energy level 2
4. Show an electron transition from energy level 2 to
3
5. Write formula for calculating this energy change
and calculate energy.
6. Give each other high fives!!
14
IBM – Almaden:
“Stadium Corral”
This image shows a ring of 76 iron atoms on a copper (111) surface. Electrons on
this surface form a two-dimensional electron gas and scatter from the iron atoms
but are confined by boundary or "corral." The wave pattern in the interior is due to
the density distribution of the trapped electrons.Their energies and spatial
distribution can be quite accurately calculated by solving the classic problem of a
quantum mechanical particle in a hard-walled box. Quantum corrals provide us with
a unique opportunity to study and visualize the quantum behavior of electrons
Heisenberg’s Uncertainty Principle (1927)
It is impossible to determine the
exact position and exact
momentum (p) of an electron.
p=mv
 To determine the position of an e-, you
have to detect how light reflects off it.
 But light means photons, which means
energy. When photons strike an e-, they
Werner
may change its motionHeisenberg
(its momentum).
(1901 – 1976)
16
Electron density distribution in H atom
17
6.5: Quantum Mechanics & Atomic Orbitals
Schrödinger’s wave function:
 Relates probability (Y2) of predicting
position of e- to its energy.
h2 d 2Y
dY
E
 UY  ih
2
2m dx
dt
Erwin
Schrödinger
(1887 – 1961)
Where:
U = potential energy
x = position
t = time
m = mass
i =√(-1)
http://daugerresearch.com/orbitals/index.shtml
18
Probability plots of 1s, 2s, and 3s orbitals
19
6.6: Representations of Orbitals
www.orbitals.com; animation 1, Draft of a letter from Bohr to Heisenberg (never
sent)
s orbital
p orbitals
20
d orbitals
f orbitals: very complicated
6.7: Filling Order of Orbitals
1. Aufbau principle: e- enter orbitals of lowest energy
first
7p
7s
6s
5s
6p
5p
4p
6d
5d
5f
x7
4f
x7
4d
3d
4s
3p
3s
2p
2s
1s
• Relative stability & average distance of e- from nucleus
22
Animation for filling of Orbitals
Use the “diagonal rule”
(some exceptions do occur).
Sub-level maxima:
s = 2 ep = 6 ed = 10 ef = 14 e…
1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d
7s 7p
23
2. Pauli exclusion principle (1925): no two e- can have
the same four quantum numbers; e- in same
orbital have opposite spins (up and down)
Wolfgang
Pauli
(1900 – 1958)
3. Hund’s rule: e- are added singly to each equivalent
(degenerate) orbital before pairing
Ex: Phosphorus (15 e-) has unpaired e- in
the valence (outer) shell.
1s 2s 2p
3s 3p
Friedrich
Hund
(1896 - 1997)
24
6.9: Periodic Table & Electronic Configurations
s block
f block
d block
p block
s2
s1 s2
1s
2s
3s
4s
5s
6s
7s
p1 p2 p3 p4 p5 p 6
d1
3d
4d f1 f2 f3 f4 f5 f6 f7 f8 f9 f10
5d
6d 4f
5f
2p
d2 d3 d5 d5 d6 d7 d8 d10d103p
4p
3d
5p
11
12
13
14
4d
f f f f
6p
5d
7p
6d
Notable Exceptions:
Cr & Mo: [Ar] 4s1 3d5 not [Ar] 4s2 3d4
Cu, Ag, & Au: [Ar] 4s13d10 not [Ar] 4s23d9
Electronic Configurations
Element
Standard Configuration
Nitrogen
1s22s22p3
Scandium
1s22s22p63s23p64s23d1
Gallium
1s22s22p63s23p64s23d104p1
Noble Gas
Shorthand
[He] 2s22p3
[Ar] 4s23d1
[Ar] 4s23d104p1
26
Noble Gas
Shorthand
Element
Standard Configuration
Lanthanum
1s2 2s22p6 3s23p6 4s23d104p6
5s24d105p6 6s25d1
[Xe] 6s25d1
Cerium
1s2 2s22p6 3s23p6 4s23d104p6
5s24d105p6 6s25d14f1
[Xe] 6s25d14f1`
1s2 2s22p6 3s23p6 4s23d104p6
Praseodymium
5s24d105p6 6s24f3
[Xe] 6s24f3
27
Electron Configuration for Ions
Valence Electrons: Only s and p e are valence
electrons. The maximum number of valence e that an
atom can have is 8. WHY? Write the electron
configurations for the following ions:
Cr +
Cr3+
Ground State Electron Config. V. Excited State
Electron Configuration
28
Ways to Represent Electron Configuration
1.Expanded Electron Configuration
2.Condensed Electron Configurations
3.Orbital Notation
4.Electron Dot Structure
Write the above four electron configurations for Zinc, Zinc ion and Cu ion.
Paramagnetic
Diamagnetic
Why are some ions colored and some aren’t?
29
Electron Configuration and
Para- and Diamagnetism
demo + activity
30