3-atomic structure overview characteristics of atoms interaction b/tw matter and light...
TRANSCRIPT
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3-Atomic Structure
Overview
• Characteristics of Atoms
• Interaction b/tw matter and light– Photoelectric Effect
• Absorption and Emission Spectra
• Electron behavior
• Quantum numbers
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Atomic Structure
• Atomic orbitals– Orbital energies– Electron configuration and the periodic
table
• Periodic table– Periodic properties– Energy
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Characteristics of Atoms
• Atoms possess mass• Atoms contain positive nuclei• Atoms contain electrons• Atoms occupy volume• Atoms have various properties• Atoms attract one another• Atoms can combine with one another to form
molecules
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Atomic Structure
• Atomic structure studied through atomic interaction with light
• Light: electromagnetic radiation– carries energy through space– moves at 3.00 x 108 m/s in vacuum– wavelike characteristics
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Electromagnetic Spectrum
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Visible Spectrum
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Wavelength () & Frequency ()
= number of complete cycles to pass given point in 1 second
amplitude
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Energy
c = x = 3.00 x 108 m/s
long wavelength low frequency
short wavelength high frequency
Low Energy
High Energy
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Energy
Mathematical relationship:
E = hE = energy
h = Planck’s constant: 6.63 x 10–34 J s
= frequency in s–1
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Energy
Mathematical relationship:E = hc = x
hc
E =
Energy: directly proportional to frequencyinversely proportional to wavelength
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Problems 3-1, 2, & 3
1. a) Calculate the wavelength of light with a frequency = 5.77 x 1014 s–1 b) What is the energy of this light?
2. Which is higher in energy, light of wave-length of 250 nm or light of 5.4 x 10–7 m?
3. a) What is the frequency of light with an energy of 3.4 x 10–19 J?b) What is the wavelength of light with an energy of 1.4 x 10–20 J?
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Photoelectric Effect
• Light on metal surface
• Electrons emitted
• Threshold frequency, o
If < o, no photoelectric effect
If > o, photoelectric effect
As , kinetic energy of electrons
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Photoelectric Effect
Einstein: energy frequency
If < o electron doesn’t have enough energy to leave the atom
If > o electron does have enough energy to leave the atom
Energy is transferred from light to electron, extra is kinetic energy of electron
Ephoton = hphoton = ho + KEelectron
KEelectron = hphoton – ho
Animation
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Problem 3-4
A given metal has a photoelectric threshold frequency of o = 1.3 x 1014
s1. If light of = 455 nm is used to produce the photoelectric effect, determine the kinetic energy of the electrons that are produced.
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Bohr Model
Line spectra
Light through a prism continuous spectrum:
Ordinary white light
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Bohr Model
Line spectra
Light from gas-discharge tube
through a prism line spectrum:
H2 discharge
tube
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Line Spectra (emission)
White light
H
He
Ne
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Line Spectra (absorption)
Light source
Gas-filled tube
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Bohr Model
For hydrogen:
22 n
1
2
1 C C = 3.29 x 1015 s–1
Niels Bohr: Electron energy in the atom is quantized.
2n n
1RE H n = 1, 2, 3,….
RH = 2.18 x 10–18 J
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Bohr ModelEatom = Eelectron = h
E = Ef – Ei = h
2
18
2Hn n
10 x 18.2
n
1RE
22
fi n
1
n
1
h
R
h
E H
Line spectrum
Photoelectric effect: n
Minus sign: free electron has zero energy
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Bohr Energy Levels
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Electrons
• All electrons have same charge and mass
• Electrons have properties of waves and particles (De Broglie)
mu
hparticle
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Heisenberg Uncertainty Principle
Cannot simultaneously know the position and momentum of electron
x = h
Recognition that classical mechanics don’t work at atomic level.
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Schrödinger Equation
Erwin Schrödinger 1926
Wave functions with discrete energies
Less empirical, more theoretical
n En
n wave functions or orbitals
n2
probability density functions
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Quantum Numbers
Each orbital defined by 3 quantum numbers
Quantum number: number that labels state of electron and specifies the value of a property
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Quantum Numbers
Principal quantum number, n (shell)
Specifies energy of electron (analogous to Bohr’s n)
Average distance from nucleus
n = 1, 2, 3, 4…..
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Quantum Numbers
Azimuthal quantum number, (subshell) = 0, 1, 2… n–1
n = 1, = 0
n = 2, = 0 or 1
n = 3, = 0, 1, or 2
Etc. 0 1 2 3 4
s p d f g
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Quantum Numbers
Magnetic quantum number, m
Describes the orientation of orbital in space
m = –….+
If = 2, m = –2, –1, 0, +1, +2
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Problem 3-5
Fill in the quantum numbers in the table below.
n m
3 0 0 3s
2 –2, –1, 0, 1, 2
0
2p
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Schrödinger Equation
Wave equations: Each electron has & E associated w/ it
Probability Density Functions: 2
-graphical depiction of high probability of finding electron
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Probability Density Functions
energy
2 probability density function
s, p, d, f, g
1s 2s
3s
Node: area of 0 electron density
Link to Ron Rinehart’s page
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Probability Density Functions
2p
Node: area of 0 electron density
3p
nodes
Link to Ron Rinehart’s page
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Electrons and Orbitals
Pauli Exclusion Principle: no two electrons in the same atom may have the same quantum numbers
Electron spin quantum number ms = ½
Electrons are spin paired within a given orbital
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Electrons and Orbitals
n = 1
= 0, m = 0, ms = ½
2 electrons possible:
1,0,0,+½ and 1,0,0,–½
2 electrons per orbital
1s1 H
1s2 He
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Electrons and Orbitals
n = 2
= 0, m = 0, ms = ½2,0,0, ½2 electrons possible
n = 2
= 1, m = –1,0,+1, ms = ½2,1,–1, ½ 2,1,0, ½ 2,1,+1, ½6 electrons possible
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Electron Configurations
n = 1
1s 2 electrons possible
H 1e– 1s1
He 2e– 1s2
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Electron Configurations
n = 2
2s 2 electrons possible
Li 3e– 1s2 2s1
1s
2s
Be 4e– 1s2 2s2
1s
2s
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Electron Configurations
n = 2
2p = 1, m = –1, 0, +1
3 x 2p orbitals (px, py, pz): 6 electrons possible
B 5e– 1s2 2s2 2p1
1s
2s
2p
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Electron Configurations
n = 2
2p = 1, m = –1, 0, +1
3 x 2p orbitals (px, py, pz): 6 electrons possible
B 5e– 1s2 2s2 2p1
1s
2s
2p
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Electron Configurations
n = 2
2p = 1, m = –1, 0, +1
C 6e– 1s2 2s2 2p2
1s
2s
2p
Hund’s Rule: for degenerate orbitals, the lowest energy is attained when electrons w/ same spin is maximized
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Problem 3-6
Write electron configurations and depict the electrons for N, O, F, and Ne.
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Electron Configurations
n = 3
3s, 3p, 3d
Na 11e– 1s2 2s2 2p63s1
1s
2s
2p
3s
3p
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Electron Configurations
n = 3
3s, 3p, 3d
Mg 12e– 1s2 2s2 2p63s2
1s
2s
2p
3s
3p
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Electron Configurations
n = 3
3s, 3p, 3d
Al 13e– 1s2 2s2 2p63s23p1
1s
2s
2p
3s
3p
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Electron Configurations
n = 3
3s, 3p, 3d
Si 14e– 1s2 2s2 2p63s23p2
1s
2s
2p
3s
3p
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Electron Configurations
n = 3
3s, 3p, 3d
P 15e– 1s2 2s2 2p63s23p3
1s
2s
2p
3s
3p
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Electron Configurations
n = 3
3s, 3p, 3d
S 16e– 1s2 2s2 2p63s23p4
1s
2s
2p
3s
3p
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Electron Configurations
n = 3
3s, 3p, 3d
Cl 17e– 1s2 2s2 2p63s23p5
1s
2s
2p
3s
3p
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Electron Configurations
n = 3
3s, 3p, 3d
Ar 18e– 1s2 2s2 2p63s23p6
1s
2s
2p
3s
3p
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Electron Configurations
3d vs. 4s
Filling order 1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f 5g
6s 6p 6d
7s 7p
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Electron Configurations
1s
2s2p
3s
3p
4s
4p
3dK
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Electron Configurations
1s
2s2p
3s
3p
4s
4p
3dCa
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Electron Configurations
1s
2s2p
3s
3p
4s
4p
3dSc
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Electron Configurations
1s
2s2p
3s
3p
4s
4p
3dTi
Link to OSU site
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Problem 3-7
Write the electron configurations for the transition metals V – Zn. Fill in the corresponding boxes to denote the electronic spin.