1. Unit 6: Chapters 11-12. Pages 295-366 ATOMIC ELECTRON CONFIGURATIONS AND PERIODICITY
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Unit 6- Atomic Electon Configurations and Periodicity

2. Bohr Model First model of the electron behavior
Vital to understanding the atom
Does not work for atoms with
more than 1 electron
Unit 6- Atomic Electon Configurations and Periodicity

3. ? Collision of Ideas Matter Light Dalton Thompson Rutherford Bohr
De Broglie
Einstein
Plank
Maxwell
Light
Newton
Unit 6- Atomic Electon Configurations and Periodicity

4. The Photoelectric Effect
Duality of Light
Wave behavior
Particle behavior
1905
Unit 6- Atomic Electon Configurations and Periodicity

5. de Broglie’s Novel Notion
1923
de Broglie’s Novel Notion
Light was “known” (thought) to be a wave, but
Einstein showed that it also acts particle-like
Electrons were particles with known mass & charge
What if ……
                       
electrons behaved as waves also
Unit 6- Atomic Electon Configurations and Periodicity

6. Evidence for de Broglie’s Notion
Diffraction pattern obtained with firing a beam of electrons through a crystal.
This can only be explained if the electron behaves as a wave!
Nobel Prize for de Broglie in 1929
Unit 6- Atomic Electon Configurations and Periodicity

7. Electron Characteristics
Extremely small mass
Located outside the nucleus
Moving at very high speeds
Have specific energy levels
Standing wave behavior
Unit 6- Atomic Electon Configurations and Periodicity

8. Baseball vs Electron
A baseball behaves as a particle and follows a predictable path.
BUT
An electron behaves as a wave, and its path cannot be predicted.
All we can do is to calculate the probability of the electron following a specific path.
Unit 6- Atomic Electon Configurations and Periodicity

9. What if a baseball behaved like an electron?
Characteristic wavelength
baseball  m
electron  0.1 nm
All we can predict is…..
Unit 6- Atomic Electon Configurations and Periodicity

10. Werner Heisenberg (1901-1976) The Uncertainty Principle
speed
position
Proposed that the dual nature of the electron places limitation on how precisely we can know both the exact location and speed of the electron
Instead, we can only describe electron behavior in terms of probability.

11. Erwin Schrodinger (1887-1961) Wave Equation & Wave Mechanics
In 1926, Austrian physicist, proposed an equation that incorporates both the wave and particle behavior of the electron
When applied to hydrogen’s 1 electron atom, solutions provide the most probable location of finding the electron in the first energy level
Can be applied to more complex atoms too!

12. Solutions to Schrodinger’s Wave Equation
Gives the most probable location of electron in 3-D space around nucleus (probability map)
- most probable
location called an
orbital
- orbitals can hold a
maximum of 2 e-

13. “Most Successful Theory of 20th Century”
Matter
Dalton
Thompson
Rutherford
Quantum
Mechanics
Bohr
De Broglie
Heisenberg
Einstein
Schrödinger
Plank
Maxwell
Wave
Mechanics
Light
Newton
Unit 6- Atomic Electon Configurations and Periodicity

14. Mathematics of waves to define orbitals
Quantum Mechanics Model Describes the arrangement and space occupied by electrons in atoms
Electron’s energy is quantized
Quantum
Mechanics
Mathematics of waves to define orbitals
(wave mechanics)
Unit 6- Atomic Electon Configurations and Periodicity

15. Bohr Model v. Quantum Mechanics
Bohr Q. Mech.
Energy
Electron
Position/Path
Unit 6- Atomic Electon Configurations and Periodicity

16. Dartboard Analogy
Suppose the size of the probability distribution is defined
as where there is a % chance of all hits being confined.
Unit 6- Atomic Electon Configurations and Periodicity

17. Quantum Mechanics Model
The electron's movement cannot be known precisely.
We can only map the probability of finding the electron at various locations outside the nucleus.
The probability map is called an orbital.
The orbital is calculated to confine 90% of electron’s range.
Unit 6- Atomic Electon Configurations and Periodicity

18. Arrangement of Electrons in Atoms
Electrons in atoms are arranged as
SHELLS (n) = distance from nucleus
1, 2, 3, …
SUBSHELLS (l) = shape of region of probability
s, p, d, f
ORBITALS (ml) = orientation in space

19. Arrangement of Electrons in Atoms
There is a relationship between the quantum number (n) and its the number of subshells.
Principal quantum number (n) = number of subshells
Unit 6- Atomic Electon Configurations and Periodicity

20. Representing s Orbitals
Unit 6- Atomic Electon Configurations and Periodicity

21. Comparison of 1s and 2s Orbitals
The 2s orbital is similar to the 1s orbital, but larger in size.
”Larger” means that the highest probability for finding the electron lies farther out from the nucleus.
Each can hold a maximum of
electrons.
Unit 6- Atomic Electon Configurations and Periodicity

22. Probability Maps of the Three 2p Orbitals
The 2p orbital is in the n = energy level.
There are p orbitals oriented in three directions.
Each orbital can hold a maximum of electrons.
The maximum number of electrons in the 2p sublevel is
Adding all 2p orbitals would result in a sphere.
Unit 6- Atomic Electon Configurations and Periodicity

23. Probability Maps of the Five 3d Orbitals
The five 3d orbitals are generally oriented in different directions.
Adding all five orbitals, would result in a sphere.
The five orbitals, taken together, make up the d subshell of the n = 3 shell.
Each orbital can hold a maximum of two electrons.
This sublevel has a maximum of electrons.
Unit 6- Atomic Electon Configurations and Periodicity

24. Probability Maps of 7 f Orbitals

25. Arrangement of Electrons in Atoms Electron Spin Quantum Number- ms
Each orbital can be assigned no more than 2 electrons! And each electron spins in opposite directions.

26. Electron Spin Quantum Number
Diamagnetic: NOT attracted to a magnetic field
Paramagnetic: substance is attracted to a magnetic field. Substance has unpaired electrons.

27. 4 QUANTUM NUMBERS Summary: n ---> shell 1, 2, 3, 4, ...
l ---> sublevel s, p, d, f
ml ---> orbital -l l
ms ---> electron spin +1/2 and -1/2

28. Pauli Exclusion Principle- No two electrons in the same atom can have the same set of 4 quantum numbers.
Determine the quantum numbers for the outer two valence electrons in the lithium atom.

29. Aufbau Principle-Electrons fill open lower energy levels sequentially lower energy to higher energy

30. Writing Electron Configurations
Two ways of writing configs. One is called the spdf notation. 1 s
value of n
value of l
no. of
electrons
spdf notation
for H, atomic number = 1

31. Broad Periodic Table Classifications
Representative Elements (main group): filling s and p orbitals (Na, Al, Ne, O)
Transition Elements: filling d orbitals (Fe, Co, Ni)
Lanthanide and Actinide Series (inner transition elements): filling 4f and 5f orbitals (Eu, Am, Es)

32. Writing Orbital Notations
Two ways of writing configs. Other is called the orbital box notation.
One electron has n = 1, l = 0, ml = 0, ms = + 1/2
Other electron has n = 1, l = 0, ml = 0, ms = - 1/2

33. Energy ordering of orbitals for multi-electron atoms
Different subshells within the same principal shell have different energies.
The more complex the subshell, the higher its energy. This explains why the 3d subshell is higher in energy than the 4s subshell.

34. Rules for Filling Orbitals
Bottom-up (Aufbau’s principle)
Fill orbitals singly before doubling up (Hund’s Rule)
Paired electrons have opposite spin (Pauli exclusion principle)
Unit 6- Atomic Electon Configurations and Periodicity

35. Valence Configuration
Cobalt
Symbol
Atomic Number
Full Configuration
Valence Configuration
Shorthand Configuration
Unit 6- Atomic Electon Configurations and Periodicity

36. Orbital diagram and electron configuration for a ground state lithium atom

37. Orbital diagram and electron configuration for a ground state carbon atom
Hund’s Rule- electrons in the same sublevel will spread out into their own orbital before doubling up.

38. Silicon's valence electrons

39. Selenium's valence electrons

40. Core electrons and valence electrons in germanium

41. Outer electron configuration for the elements

42. The periodic table gives the electron configuration for As

43. Valence Electrons by Group

44. Ion charges by group

45. Periodic Law Example: Group 2 Be 2, 2 Mg 2, 8, 2 Ca 2, 2, 8, 2
All the elements in a group have the same electron configuration in their outermost shells
Example: Group 2
Be 2, 2
Mg 2, 8, 2
Ca 2, 2, 8, 2

46. Question A. Cl and Br B. P and S C. O and S
Specify if each pair has chemical properties that are similar (1) or not similar (2):
A. Cl and Br
B. P and S
C. O and S

47. General Periodic Trends
1. Atomic and ionic size 2. Electron affinity
3. Ionization energy Metallic Character
Higher effective nuclear charge
Electrons held more tightly
Larger orbitals.
Electrons held less
tightly.

48. Effective Nuclear Charge, Z*
Z* is the nuclear charge experienced by the outermost electrons. Screen 8.6.
Explains why E(2s) < E(2p)
Z* increases across a period owing to incomplete shielding by inner electrons.
Estimate Z* by --> [ Z - (no. inner electrons) ]
Z = number of electrons
Charge felt by 2s e- in Li Z* = = 1
Be Z* = = 2
B Z* = = 3 and so on!

49. Effective Nuclear Charge
Figure 8.6
Electron cloud for 1s electrons

50. Effective Nuclear Charge, Z*
Atom Z* Experienced by Electrons in Valence Orbitals Li Be B C N O F
Increase in Z* across a period

51. Beryllium
Lithium
Sodium

52. Atomic Size Size goes UP on going down a group. See Figure 8.9.
Because electrons are added further from the nucleus, there is less attraction.
Size goes DOWN on going across a period.

53. Atomic Radii
Figure 8.9

54. Trends in Atomic Size See Figures 8.9 & 8.10

55. Ion Sizes Does the size go
up or down when losing an electron to form a cation?

56. Ion Sizes CATIONS are SMALLER than the atoms from which they come.Li +
, 78 pm
2e and 3 p
Forming a cation.
Li,152 pm
3e and 3p
CATIONS are SMALLER than the atoms from which they come.

57. Ion Sizes
Does the size go up or down when gaining an electron to form an anion?

58. Ion Sizes ANIONS are LARGER than the atoms from which they come.-
, 133 pm
10 e and 9 p
F, 71 pm
9e and 9p
Forming an anion.
ANIONS are LARGER than the atoms from which they come.

59. Trends in Ion Sizes
Figure 8.13

60. Ionization Energy See Screen 8.12
IE = energy required to remove an electron from an atom in the gas phase.
Mg (g) kJ ---> Mg+ (g) + e-

61. Ionization Energy See Screen 8.12
Mg (g) kJ ---> Mg+ (g) + e-
Mg+ (g) kJ ---> Mg2+ (g) + e-
Mg2+ (g) kJ ---> Mg3+ (g) + e-
Energy cost is very high to dip into a shell of lower n. This is why ox. no. = Group no.

62. Trends in Ionization Energy

63. Trends in Ionization Energy
IE increases across a period because Z* increases.
Metals lose electrons more easily than nonmetals.
Metals are good reducing agents.
Nonmetals lose electrons with difficulty.

64. Trends in Ionization Energy
IE decreases down a group
Because size increases.
Reducing ability generally increases down the periodic table.
See reactions of Li, Na, K

65. Electronegativity
A measure of the ability of an atom that is bonded to another atom to attract electrons to itself.
                                                                                                             
                                                                                                             

66. Electron Affinity A few elements GAIN electrons to form anions.
Electron affinity is the energy involved when an atom gains an electron to form an anion.
A(g) + e- ---> A-(g)
E.A. = ∆E

67. Electron Affinity of Oxygen
∆E is EXOthermic because O has an affinity for an e-.
[He]  
O atom
+ electron O
[He]   -
ion
EA = kJ

68. Trends in Electron Affinity
See Figure 8.12 and Appendix F
Affinity for electron increases across a period (EA becomes more negative).
Affinity decreases down a group (EA becomes less negative).
Atom EA
F -328 kJ
Cl -349 kJ
Br -325 kJ
I -295 kJ

69. Trends in Electron Affinity

70. Metallic character trends in the periodic table

71. Metallic Character
The text links metallic character to the tendency to lose electrons in chemical reactions, and nonmetallic character to the tendency to gain electrons in chemical reactions. The metallic character trends therefore follow the ionization energy trends

72. The metallic character trends explain the location of metals, metalloids, and nonmetals

73. Which is the more metallic element, Sn or Te?

74. Which is the more metallic element, Si or Sn?