1. Chapter 4: Arrangement of Electrons in Atoms
Chemistry

2. Development of a New Atomic Model
There were some problems with the Rutherford model…It did not answer:
Where the e- were located in the space outside the nucleus
Why the e- did not crash into the nucleus
Why atoms produce spectra (colors) at specific wavelengths when energy is added

3. Properties of Light Wave-Particle Nature of Light – early 1900’s
A Dual Nature
It was discovered that light and e- both have wave-like and particle-like properties

4. Wave Nature of Light
Electromagnetic radiation – form of energy that exhibits wave-like behavior as it travels through space
Electromagnetic spectrum
All the forms of electromagnetic radiation
Speed of light in a vacuum
3.0 x 108 m/s

5. Wave Nature of Light Wavelength Frequency
Distance between two corresponding points on adjacent waves λ nm
Frequency
Number of waves that pass a given point in a specified time -u
Hz - Hertz

6. Wave Nature of Light Figure 4-1, page 92 Equation Spectroscope c=λu
Speed = wavelength * frequency
Indirectly related!
Spectroscope
Device that separates
light into a spectrum that
can be seen

7. Particle Nature of Light
Quantum
Minimum quantity of energy that can be lost or gained by an atom
Equation
E=hu
Direct relationship between quanta (particle nature) and frequency (wave nature)
Planck’s Constant (h)
h=6.626 x Js

8. Particle Nature of Light
Photon
Individual quantum of light; “packet”
The Hydrogen Atom
Line emission spectrum (Figure 4-5, page 95)
Ground State
Lowest energy state (closest to the nucleus)
Excited State
State of higher energy
Each element has a characteristic bright-line spectrum – much like a fingerprint!**

10. Particle Nature of Light
Why does an emission spectrum occur?
Atoms get extra energy – ex. voltage – and the e- jumps from ground state to excited state
Atoms return to original energy, e- drops back down to ground state
The energy is transferred out of the atom in a NEW FORM
Continuous spectrum
Emission of continuous range of frequencies
Line Emission Spectrum
Shows distinct lines

11. Bohr Model of the Hydrogen Atom
Described electrons as PARTICLES
1913 – Danish physicist – Niels Bohr
Single e- circled around nucleus in allowed paths or orbits
e- has fixed E when in this orbit (lowest E closest to nucleus)
Lot of empty space between nucleus and e- in which e- cannot be in
E increases as e- moves to farther orbits

12. Bohr Model (cont) ONLY explained atoms with one e-
Therefore – only worked with hydrogen!!
The principles of his work is applied to the models of other atoms, but the models do not perfectly fit the experimental data.

13. Orbits = The circular paths electrons followed in the Bohr model of the atom
Spectroscopy
Study of light emitted by excited atoms
Bright line spectrum

14. The Quantum Model of the Atom
e- act as both waves and particles!!
De Broglie
1924 – French physicist
e- may have a wave-particle nature
Would explain why e- only had certain orbits
Diffraction
Bending of wave as it passes by edge of object
Interference
Occurs when waves overlap

15. The Quantum Model of the Atom
Heisenberg Uncertainty Principle
1927 – German physicist
It is impossible to determine simultaneously both the position and velocity of an e-
12:28-14:28

17. The Quantum Model of the Atom
Schrodinger Wave Equation
1926 – Austrian physicist
Applies to all atoms, treats e- as waves
Nucleus is surrounded by orbitals
Laid foundation for modern quantum theory
Orbital – 3D region around nucleus in which an e- can be found
Cannot pinpoint e- location!!

18. Quantum Numbers Quantum Numbers Solutions to Schrodinger’s wave eqn
Probability of finding an e-
“address” of e-
Four Quantum Numbers
Principle
Angular Momentum
Magnetic
Spin

19. Principle Quantum Number
Which main energy level? (“shell”)
The distance from the nucleus
Symbol- n
n is normally 1-7
Greater n value means farther from the nucleus

20. Angular Momentum Quantum Number
What is the shape of the orbital?
Symbol – l
l = s,p,d,f

21. Magnetic Quantum Number
Orientation of orbital around nucleus
Symbol – ml
s – 1
p – 3
d – 5
f – 7
Every orientation can hold 2 e-!!
A “subshell” is made of all of the orientations of a particular shape of orbital
Figures 4-13, 4-14, 4-15 on page

22. Spin Quantum Number Each e- in one orbital must have opposite spins
Symbol – ms
+ ½ , - ½
Two “allowed” values and corresponds to direction of spin

23. Electron Configuration
Electron configurations – arrangements of e- in atoms
Rules:
Aufbau Principle – an e- occupies the lowest energy first
Hund’s Rule – place one electron in each equal energy orbital before pairing
Pauli Exclusion Principle – no 2 e- in the same atom can have the same set of QN
14:30-18:25

24. Electron Configuration
Representing electron configurations
Use the periodic table to write!
Know the s,p,d,f block and then let your fingers do the walking!

25. Electron Configuration
Lags 1 behind
Lags 2 behind

26. Representing Electron Configurations
Three Notations
Orbital Notation
Electron Configuration Notation
Electron Dot Notation

27. Orbital Notation
Uses a series of lines and arrows to represent electrons
Examples

28. Orbital Notation
More examples

29. Electron Configuration Notation
Long Form: Eliminates lines and arrows; adds superscripts to sublevels to represent electrons
Long form examples

30. Electron Configuration Notation
Short form examples – “noble gas configuration”

31. Electron Dot Notation
Outer shell e- - Outermost electrons; In highest principle quantum #
Inner shell e- - not in the highest energy level
Highest occupied energy level / highest principle quantum number
Valence electrons – outermost e-
Examples

32. Electron Dot Notation
More examples

33. Summary Questions How many orbitals are in a d subshell?
How many individual orbitals are found in Principle Quantum #3 (the third main energy level)
How many orbital shapes are found in Principle Quantum #2?
How many electrons can be found in the fourth energy level?
A single 4s orbital can hold how many electrons?

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