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Electronic Structure of Atoms (i.e., Quantum Mechanics)

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Electronic Structure of Atoms (i.e., Quantum Mechanics) Brown, LeMay Ch 6 AP Chemistry 6.1: Light is a Wave Electromagnetic spectrum: A form of radiant energy (can ... – PowerPoint PPT presentation

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Title: Electronic Structure of Atoms (i.e., Quantum Mechanics)


1
Electronic Structure of Atoms(i.e., Quantum
Mechanics)
  • Brown, LeMay Ch 6
  • AP Chemistry

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)
  • Speed of Light c 3.0 x 108 m/s (in a vacuum)
  • Wavelength (l) distance between wave peaks
    (determines color of light)
  • Frequency (n) cycles/sec (measured in Hz)

c l n
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 lt white hot lt blue hot
  • Plancks constant
  • Blackbody radiation can be explained if energy
    can be released or absorbed in packets of a
    standard size he called quanta (singular
    quantum).
  • where Plancks constant (h) 6.63 x 10-34 J-s

Max Planck(1858-1947)
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.

Albert Einstein(1879-1955)
5
6.3 Bohrs Model of the H Atom (and only H!)
  • Atomic emission spectra
  • Most sources produce light that contains many
    wavelengths at once.
  • 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).
  • Atomic emission spectra are inverses of atomic
    absorption spectra.

6
  • 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
  • principal quantum number (n) 1, 2, 3, 4,
  • Rydbergs constant (RH) 2.178 x 10-18 J

Niels Bohr(1888-1962)
Johannes Rydberg(1854-1919)
7
5 4 3 2 1
E5 E4 E3 E2 E1
Increasing Energy, E
Principal Quantum Number, n
  • As n approaches 8, the e- is essentially removed
    from the atom, and E8 0.
  • ground state lowest energy level in which an e-
    is stable
  • excited state any energy level higher than an
    e-s ground state

8
  • ni initial orbital of e-
  • nf final orbital of e- in its transition

9
Theodore Lyman (1874 - 1954)
5 4 3 2 1
FriedrichPaschen(1865 - 1947)
n
?
JohannBalmer(1825 1898)
FrederickBrackett(1896 1988)
Figure 1 Line series are transitions from one level to another. Figure 1 Line series are transitions from one level to another. 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 2 Visible
Paschen 3 IR
Brackett 4 Far IR
10
6.4 Matter is a Wave
  • Planck said E h c / l
  • Einstein said E m c2
  • Louis DeBroglie said (1924) h c / l m c2
  • h / l m c
  • Therefore

Louisde Broglie(1892 - 1987)
m h / cl Particles (with mass) have an associated wavelength
l h / mc Waves (with a wavelength) have an associated mass and velocity
11
IBM AlmadenStadium 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 within small confining structures.

12
Heisenbergs Uncertainty Principle (1927)
  • It is impossible to determine the exact position
    and exact momentum (p) of an electron.
  • p m v
  • 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 may change its
    motion (its momentum).

WernerHeisenberg(1901 1976)
13
Electron density distribution in H atom
14
6.5 Quantum Mechanics Atomic Orbitals
  • Schrödingers wave function
  • Relates probability (Y2) of predicting position
    of e- to its energy.

ErwinSchrödinger(1887 1961)
  • Where U potential energy
  • x position t time
  • m mass i v(-1)

15
Probability plots of 1s, 2s, and 3s orbitals
16
6.6 Representations of Orbitals
  • s orbital
  • p orbitals

17
  • d orbitals
  • f orbitals very complicated

18
Figure 2 Orbital Quantum Numbers
Symbol Name Description Meaning Equations
n Principle Q.N. Energy level (i.e. Bohrs theory) Shell number n 1, 2, 3, 4, 5, 6, 7 n 1, 2, 3,
l Angular Momentum Q.N. General probability plot (shape of the orbitals) Subshell number l 0, 1, 2, 3   l 0 means s l 1 means p l 2 means d l 3 means f l 0, 1, 2, , n 1   Ex If n 1, l can only be 0 if n 2, l can be 0 or 1.
19
Symbol Name Description Meaning Equations
ml Magnetic Q.N. 3-D orientation of the orbital s has 1 p has 3 d has 5 f has 7 ml -l, -l 1, , 0, l, , l   There are (2l 1) values.  
ms Spin Q.N. Spin of the electron Parallel or antiparallel to field ms ½ or -½
s, p, d, and f come from the words sharp,
principal, diffuse, and fundamental.
20
Permissible Quantum Numbers
  • (4, 1, 2, ½)
  • (5, 2, 0, 0)
  • (2, 2, 1, ½)

Not permissible if l 1, ml 1, 0, or 1 (p
orbitals only have 3 subshells)
Not permissible ms ½ or ½
Not permissible if n 2, l 0 or 1 (there is
no 2d orbital)
21
6.7 Filling Order of Orbitals
  1. Aufbau principle e- enter orbitals of lowest
    energy first ( postulated by Bohr, 1920)

7p
6d
6p
5d
5p
4d
4p
3d
3p
2p
  • Relative stability average distance of e- from
    nucleus

22
6.7 Filling Order of Orbitals
  1. Aufbau principle e- enter orbitals of lowest
    energy first
  • Relative stability average distance of e- from
    nucleus

23
1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p
5d 5f 6s 6p 6d 7s 7p
  • Use the diagonal rule
  • (some exceptions do occur).
  • Sub-level maxima
  • s 2 e-
  • p 6 e-
  • d 10 e-
  • f 14 e-

24
  • Pauli exclusion principle (1925) no two e- can
    have the same four quantum numbers e- in same
    orbital have opposite spins (up and down)
  • Hunds rule e- are added singly to each
    equivalent (degenerate) orbital before pairing
  • Ex Phosphorus (15 e-) has unpaired e- inthe
    valence (outer) shell.
  • 1s 2s 2p 3s 3p

WolfgangPauli(1900 1958)
FriedrichHund(1896 - 1997)
25
6.9 Periodic Table Electronic Configurations
s block
p block
d block
f block
s2
s1
s2
1s 2s 3s 4s 5s 6s 7s
2p 3p 4p 5p 6p 7p
d1
3d 4d 5d 6d
3d 4d 5d 6d
4f 5f
26
Electronic Configurations
Element Standard Configuration Noble Gas Shorthand
Nitrogen
Scandium
Gallium
He 2s22p3
1s22s22p3
1s22s22p63s23p64s23d1

Ar 4s23d1
Ar 4s23d104p1
1s22s22p63s23p64s23d104p1
27
Element Standard Configuration Noble Gas Shorthand
Lanthanum
Cerium
Praseodymium
Xe 6s25d1
1s2 2s22p6 3s23p6 4s23d104p6 5s24d105p6 6s25d1
Xe 6s25d14f1

1s2 2s22p6 3s23p6 4s23d104p6 5s24d105p6 6s25d14f1
Xe 6s24f3
1s2 2s22p6 3s23p6 4s23d104p6 5s24d105p6 6s24f3
28
Notable Exceptions
  • Cr Mo Ar 4s1 3d5 not Ar 4s2 3d4
  •  
  • Cu, Ag, Au Ar 4s13d10 not Ar 4s23d9
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