Atomic Structure
26
The Wave Nature of Light • All forms of NRG/Light have characteristic wavelengths (λ) and frequency (υ). – Inversely related • λ υ = c (the speed of light) – Light visible to the naked eye exists as a tiny portion of the electromagnetic spectrum
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Transcript of Atomic Structure
The Wave Nature of Light
• All forms of NRG/Light have characteristic wavelengths (λ) and frequency (υ).– Inversely related
• λ υ = c (the speed of light)–Light visible to the naked eye exists as a
tiny portion of the electromagnetic spectrum
Electromagnetic Spectrum
Max Planck
• Transfer of energy was not continuous– Only came in certain
values (quantized)• ΔE = hν
– h = Planck’s constant = 6.626 x 10-34 Js
• Packets of energy (quantum
Albert Einstein
• Proposed that electromagnetic radiation was quantized and made up of a stream of particles– Photons
• The dual nature of light
• λ = h/mv (deBroglie equation)
Electrons as Waves
• Louis de Broglie (1924)
– Applied wave-particle theory to electrons– electrons exhibit wave properties
QUANTIZED WAVELENGTHS
Adapted from work by Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Standing Wave 200
150
100
50
0
- 50
-100
-150
-2000 50 100 150 200
Second Harmonic or First Overtone 200
150
100
50
0
- 50
-100
-150
-2000 50 100 150 200
Fundamental mode 200
150
100
50
0
- 50
-100
-150
-2000 50 100 150 200
Louis de Broglie~1924
Electrons as WavesQUANTIZED WAVELENGTHS
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Electrons as Waves
Evidence: DIFFRACTION PATTERNS
ELECTRONSVISIBLE LIGHT
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chemDavis, Frey, Sarquis, Sarquis, Modern Chemistry 2006, page 105
Spectrum
• Light is diffracted (bent) through different objects
• Hydrogen Spectra gave interesting results
Niels Bohr
• In the Bohr Model (1913) the neutrons and protons occupy a dense central region called the nucleus, and the electrons orbit the nucleus much like planets orbiting the Sun.
• They are not confined to a planar orbit like the planets are.
Bohr Model
After Rutherford’s discovery, Bohr proposed that electrons travel in definite orbits around the nucleus.
Planetary model
Bohr Cont.
• When the electrons are in their lowest possible NRG level, they are in their ground state.
• Electrons absorb NRG and go to a higher NRG level(Excited State)
• NRG (Light) is released when the electron jumps from a higher NRG level to a lower NRG level(Fluorescence)– Constantly happening
An unsatisfactory model for the hydrogen atom
According to classical physics, lightshould be emitted as the electron circles the nucleus. A loss of energywould cause the electron to be drawncloser to the nucleus and eventuallyspiral into it.
Hill, Petrucci, General Chemistry An Integrated Approach 2nd Edition, page 294
Quantum Mechanical Model
Modern atomic theory describes the electronic structure of the atom as the probability of finding electrons within certain regions of space (orbitals).
Niels Bohr &Albert Einstein
Modern View
• The atom is mostly empty space• Two regions
– Nucleus • protons and neutrons
– Electron cloud• region where you might find an electron
Quantum Mechanics
• Heisenberg Uncertainty Principle– Impossible to know both the velocity and
position of an electron at the same time
Microscope
Electron
g
Werner Heisenberg~1926
Quantum Mechanics
σ3/2 Zπ
11s 0
eΨ a
• Schrödinger Wave Equation (1926)
– finite # of solutions quantized energy levels
– defines probability of finding an electron
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Erwin Schrodinger~1926
Quantum Mechanics
• Orbital (“electron cloud”)– Region in space where there is 90% probability
of finding an electron
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Electron Probability vs. Distance
Ele
ctro
n P
roba
bilit
y (%
)
Distance from the Nucleus (pm)
100 150 200 2505000
10
20
30
40
Orbital
90% probability offinding the electron
Quantum Numbers
UPPER LEVEL
• Four Quantum Numbers:– Specify the “address” of each electron in an
atom
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Quantum Numbers
Principal Quantum Number ( n )
Angular Momentum Quantum # ( l )
Magnetic Quantum Number ( ml )
Spin Quantum Number ( ms )
Quantum Numbers
1. Principal Quantum Number ( n )
– Energy level
– Size of the orbital
– Integral values
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
1s
2s
3s
Quantum Numbers
s p d f
2. Angular Momentum Quantum # ( l )– Energy sublevel
– Shape of the orbital
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Angular Momentum Quantum # ( l )
• Has integral values from 0 to n-1– Related to the shape of the orbital
• l = 0 is called s• l = 1 is called p• l = 2 is called d• l = 3 is called f
Shapes of s, p, and d-Orbitals
Shape of f-orbitals
Quantum Numbers3. Magnetic Quantum Number ( ml )
– Orientation of orbital– Specifies the exact orbital within each sublevel– Has values between l and -l
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Quantum Numbers
4. Spin Quantum Number ( ms )– Electron spin +½ or -½
– An orbital can hold 2 electrons that spin in opposite directions.
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
