1. Chapter 4 Arrangement of Electrons in Atoms
4.1 The Development of a New Atomic Model

2. Waves: Not Just For The Beach Anymore!
Properties of Light
Electromagnetic Radiation:
EM radiation are forms of energy which move through space as waves
There are many different types of EM waves
visible light
x-rays
ultraviolet light
infrared light
radio waves
Waves: Not Just For The Beach Anymore!

3. EM Waves Move at speed of light: 3.00 x 108 m/s
Speed is equal to the frequency times the wavelength c = v
Frequency (v) is the number of waves passing a given point in one second
Wavelength () is the distance between peaks of adjacent waves
Speed of light is a constant, so v is also a constant; v and  must be inversely proportional

5. Light and Energy: The Photoelectric Effect
Electrons are emitted from a metal when light shines on the metal
Incoming EM radiation from the left ejects electrons, depicted as flying off to the right, from a substance.
Radiant energy is transferred in units (or quanta) of energy called photons
(Max Planck)

6. Photoelectric Effect absorption spectrum
energy
absorption spectrum p+
ground state e- no e-
When a specific or quantized amount of energy is exposed to the atom, the electron jumps from its “ground” or original state to an “excited” state

7. travels at the speed of light
When the “excited” electron returns to lower energy levels, it releases energy in the form of light
Photoelectric Effect
energy photon
emission spectrum! p+
excited state e- no e-
travels at the speed of light
(3.00 x 108 m/s)

8. A photon is a particle of energy having a rest mass of zero and carrying a quantum of energy
A quantum is the minimum amount of energy that can be lost or gained by an atom
Energy of a photon is directly proportional to the frequency of radiation
E = hv (h is Planck’s constant,
x J * sec)

9. Electromagnetic Spectrum
Wavelength increases→
Frequency decreases→
Energy decreases→

10. Electromagnetic Spectrum

11. Wave-Particle Duality
Energy travels through space as waves, but can be thought of as a stream of particles (Einstein)
Each particle has 1 quantum of energy.

12. Line Spectrums Ground State: The lowest energy state of an atom
Excited State: A state in which an atom has a higher potential energy than in its ground state
example: Neon lights

13. Emissions Spectrum
Bright line spectrum: Light is given off by excited atoms as they return to lower energy states
Light is given off in very definite wavelengths
A spectroscope reveals lines of particular colors- light passed through a prism; specific frequencies given off.

14. The Hydrogen Line Spectrum
Definite frequency
Definite wavelength

15. Bohr Model e-
Niels Bohr p+ no
Electrons circle around the nucleus on their energy level
Energy levels

16. The Bohr Model of the Atom
Electron Orbits, or Energy Levels
Electrons can circle the nucleus only in allowed paths or orbits
The energy of the electron is greater when it is in orbits farther from the nucleus
The atom achieves the ground state when atoms occupy the closest possible positions around the nucleus
Electromagnetic radiation is emitted when electrons move closer to the nucleus.

17. The Bohr Atomic Model

18. Energy transitions Energies of atoms are fixed and definite quantities
Energy transitions occur in jumps of discrete amounts of energy
Electrons only lose energy when they move to a lower energy state

19. Shortcomings of the Bohr Model
Doesn't work for atoms larger than hydrogen
(more than one electron)
Doesn't explain chemical behavior

20. Chapter 4 Arrangement of Electrons in Atoms
4.2 The Quantum Model of the Atom

21. Electrons as Waves and Particles
Louis deBroglie (1924)
Electrons have wavelike properties
Consider the electron as a wave confined to a space that can have only certain frequencies

22. The Heisenberg Uncertainty Principle
"It is impossible to determine simultaneously both the position and velocity of an electron or any other particle.”
Werner Heisenberg- 1927
Electrons are located by their interactions with photons
Electrons and photons have similar energies
Interaction between a photon and an electron knocks the electron off of its course

23. The SchrØdinger Wave Equation
Proved quantization of electron energies and is the basis for Quantum Theory
Quantum theory describes mathematically the wave properties
of electrons and other very small particles
Electrons do not move around the nucleus in "planetary orbits"
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24. The SchrØdinger Wave Equation
Electrons exist in regions called orbitals
An orbital is a three-dimensional region around the nucleus that indicates the probable location of an electron
SchrØdinger equation for probability of a single electron being found along a single axis (x-axis)

25. Atomic Orbitals & Quantum Numbers
Quantum Numbers specify the properties of atomic orbitals and the properties of the electrons in orbitals:
Principal Quantum Number (n)
Angular Momentum Quantum Number (l)
Magnetic Quantum Number (m)
Spin Quantum Number

26. Principal Quantum Number (n)
Indicates the main energy levels occupied by the electron
Values of n are positive integers
n = 1 is closest to the nucleus, and lowest in energy
The number of orbitals possible per energy level (or "shell") is equal to n2

27. Quantum Theory
Energy levels, n e-
n = 1 p+
n = 2 no
n = 3
n = 4

28. Angular Momentum Quantum Number (l)
Indicates the shape of the orbital
Number of orbital shapes = n
Possible values are l = 0, 1, 2, or 3
Shapes are designated s, p, d, f
S shape is spherical
P shape is a dumbbell, or figure 8
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29. Magnetic Quantum Number (m)
The orientation of the orbital around the nucleus
s orbitals have only one possible orientation,
m = 0
p orbitals have three possible,
m = +1, 0 or -1
d orbitals have five possible,
m = -2, -1, 0, +1, or +2
f orbitals have 7 possible orientations

30. Spin Quantum Number
Indicates the fundamental spin states of an electron in an orbital
Two possible values for spin, +1/2, -1/2
A single orbital can contain only two electrons, which must have opposite spins

31. Diagrams of the Orbitals

32. Summary of the Quantum Numbers

33. Chapter 4 Arrangement of Electrons in Atoms
4.3 Electron Configurations

34. Energy Levels

35. Writing Electron Configurations
Pauli Exclusion Principle
No two electrons in the same atom can have the same set of four quantum numbers
Aufbau Principle
An electron occupies the lowest-energy orbital that can receive it
Rules:
Hund's Rule
Orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron, and all electrons in singly occupied orbitals must have the same spin

36. Pauli Exclusion Principle
Orbital Notation (1 of 3)
Unoccupied orbitals are represented by a line, _____
Lines are labeled with the principal quantum number and the sublevel letter
Arrows are used to represent electrons
Arrows pointing up and down indicate opposite spins:
Pauli Exclusion Principle
No two electrons in the same atom can have the same set of four quantum numbers (occupy the same the same time)

37. Writing Electron Configurations
Hund's Rule
Orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron, and all electrons in singly occupied orbitals must have the same spin
1st e-
2nd e-
3rd e-
4th e-

38. Configuration Notation (2 of 3)
The number of electrons in a sublevel is indicated by adding a superscript to the sublevel designation
Hydrogen = 1s1
Helium = 1s2
Lithium = 1s2 2s1
Aufbau Principle
An electron occupies the lowest-energy orbital that can receive it
(Always start with n = 1 and work your way up)

39. Sublevel Blocks on the Periodic Table1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18
Sublevel Blocks on the Periodic Table

41. Fill-In the Periodic Table using Energy Levels & Sublevels

42. Orbital Filling Order START 1s 2s 2p 3p 3d 3s 4f 4p 4d 4s 5p 5d 5f 5s
FINISH

43. Noble Gas Notation (3 of 3)
The configuration begins with the preceding noble gas’s symbol in brackets and is followed by the rest of the configuration for the particular element.
[Ne] 3s23p5

44. Terms Highest occupied energy level:
The electron containing energy level with the highest principal quantum number
Inner shell electrons:
Electrons that are not in the highest energy level
Octet Rule:
Highest energy level s and p electrons are filled (8 electrons)

45. Octet Rule Characteristic of noble gases, Group 18
Exceptions: Hydrogen & Helium
Noble gas configuration:
Outer main energy level fully occupied, usually (except for He) by eight electrons
This configuration has extra stability

46. Survey of the Periodic Table
Elements of the Fourth Period
Irregularity of Chromium
Expected: 1s22s22p63s23p64s23d4
Actual: 1s22s22p63s23p64s13d5
Rule:
Sublevels are most stable when they are either half or completely filled. Electrons will shift to different energy levels to accommodate this stability whenever possible.

47. Survey of the Periodic Table
Several transition and rare-earth elements borrow from smaller sublevels in order to half fill larger sublevels
i.e. d borrows 1 e- from s
This accounts for some of the unexpected electron configurations found within the transition elements.

48. Magnetic Fields & Electrons
Paramagnetic:
When an atom has unpaired electrons, it will be attracted into a magnetic field
Ex: 1s22s22p2
Dimagnetic:
When an atom has only paired electrons, it will be slightly repelled by a magnetic field
Ex: 1s22s22p63s2