1. Atomic Theory

2. I. Waves & Particles (p. 91 - 94)
Ch. 4 - Electrons in Atoms
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3. B. EM Spectrum
HIGH
ENERGY
LOW
ENERGY
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4. B. EM Spectrum HIGH LOW ENERGY ENERGY R O Y G. B I V red orange yellow
green
blue
indigo
violet
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5. C. Quantum Theory Planck (1900)
Observed - emission of light from hot objects
Concluded - energy is emitted in small, specific amounts (quanta)
Quantum - minimum amount of energy change
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6. C. Quantum Theory Einstein (1905) Observed - photoelectric effect
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7. “wave-particle duality”
C. Quantum Theory
Einstein (1905)
Concluded - light has properties of both waves and particles
“wave-particle duality”
Photon - particle of light that carries a quantum of energy
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8. C. Quantum Theory
The energy of a photon is proportional to its frequency.
E = h
E: energy (J, joules)
h: Planck’s constant (  J·s)
: frequency (Hz)
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9. II. Bohr Model of the Atom (p. 94 - 97)
Ch. 4 - Electrons in Atoms
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10. A. Line-Emission Spectrum
excited state
ENERGY IN
PHOTON OUT
ground state
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11. B. Bohr Model
e- exist only in orbits with specific amounts of energy called energy levels
Therefore…
e- can only gain or lose certain amounts of energy
only certain photons are produced
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12. B. Bohr Model 6
Energy of photon depends on the difference in energy levels
Bohr’s calculated energies matched the IR, visible, and UV lines for the H atom 5 4 3 2 1
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13. C. Other Elements Helium
Each element has a unique bright-line emission spectrum.
“Atomic Fingerprint”
Helium
Bohr’s calculations only worked for hydrogen! 
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14. III. Quantum Model of the Atom (p. 98 - 104)
Ch. 4 - Electrons in Atoms
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15. A. Electrons as Waves QUANTIZED WAVELENGTHS Louis de Broglie (1924)
Applied wave-particle theory to e-
e- exhibit wave properties
QUANTIZED WAVELENGTHS
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16. A. Electrons as Waves
QUANTIZED WAVELENGTHS
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17. A. Electrons as Waves EVIDENCE: DIFFRACTION PATTERNS VISIBLE LIGHT
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18. B. Quantum Mechanics Heisenberg Uncertainty Principle
Impossible to know both the velocity and position of an electron at the same time
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19. B. Quantum Mechanics Schrödinger Wave Equation (1926)
finite # of solutions  quantized energy levels
defines probability of finding an e-
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20. Radial Distribution Curve
B. Quantum Mechanics
Orbital (“electron cloud”)
Region in space where there is 90% probability of finding an e-
Orbital
Radial Distribution Curve
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21. C. Quantum Numbers Four Quantum Numbers:
Specify the “address” of each electron in an atom
UPPER LEVEL
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22. C. Quantum Numbers 1. Principal Quantum Number ( n ) Energy level
Size of the orbital
n2 = # of orbitals in the energy level
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23. C. Quantum Numbers 2. Angular Momentum Quantum # ( l ) Energy sublevel
Shape of the orbital f d s p
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24. C. Quantum Numbers n = # of sublevels per level
n2 = # of orbitals per level
Sublevel sets: 1 s, 3 p, 5 d, 7 f
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25. C. Quantum Numbers 3. Magnetic Quantum Number ( ml )
Orientation of orbital
Specifies the exact orbital within each sublevel
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26. C. Quantum Numbers px py pz
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27. C. Quantum Numbers Orbitals combine to form a spherical shape. 2s 2px
2pz
2py
2px
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28. C. Quantum Numbers 4. Spin Quantum Number ( ms )
Electron spin  +½ or -½
An orbital can hold 2 electrons that spin in opposite directions.
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29. C. Quantum Numbers Pauli Exclusion Principle 1. Principal # 
No two electrons in an atom can have the same 4 quantum numbers.
Each e- has a unique “address”:
1. Principal # 
2. Ang. Mom. # 
3. Magnetic # 
4. Spin # 
energy level
sublevel (s,p,d,f)
orbital
electron
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30. Feeling overwhelmed?
Read Section 4-2!
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