1. Electronic Structure and the Periodic Table
4/21/2017
Electronic Structure and the Periodic Table
Unit 6 Honors Chemistry
Dr. Mihelcic Honors Chemistry

2. Electromagnetic Waves:
progressive, repeating disturbances that come from the movement of electric charges
Electromagnetic Waves & Light

3. Wavelength and Frequency
Wavelength (, lambda): distance between any two points in a wave
measured in any distance unit
(mainly nm or m:
1 nm = 1x10-9 m)

4. Wavelength Can be Measured in One of Two Ways…

5. Wavelength and Frequency
Frequency (; pronounced nu):
the number of cycles
of the wave that pass through a point in a unit of time
Measured in sec-1 (/sec)
1 sec-1 = 1 Hertz (Hz)

6. Illustration of Frequency

7. Wavelength is indirectly proportional to frequency
As Wavelength increases, frequency _________________.
As Wavelength decreases, frequency _________________.

8. Amplitude
Note: height of wave is amplitude (intensity or brightness of wave)
Amplitude is INDEPENDENT of frequency or wavelength!

9. Speed Speed (c): The speed of light! c = 3.00 x 108 m/s
(rounded to 3 sig figs)

10. Equation
One equation relates speed, frequency and wavelength:
c =  

11. Example
The wavelength of the radiation which produced the yellow color of sodium vapor light is nm. What is the frequency of this radiation?

12. The electromagnetic spectrum
complete range of wavelengths and frequencies
mostly invisible

13. What is color?
TED Talk: What is color?

14. The visible/continuous spectrum
continuous spectrum: components of white light split into its colors, ROY G BIV
from 390 nm (violet) to 760 nm (red)

15. Line Spectra
Pattern of lines produced by light emitted by excited atoms of an element
unique for every element
used to identify unknown elements

16. How do we see color?
TED Talk: How we see color

17. Max Planck Equation: E (Energy of a photon)= h
Light is generated as a stream of particles called PHOTONS
Equation:
E (Energy of a photon)= h
(h =Plank’s constant=
6.626x10-34Js)

18. Relationships in Planck’s Eqn.
E = h • 
High frequency, low λ, high E.
Low frequency, high λ, low E.

19. Photoelectric effect – Nobel Prize in Physics 1921 to Einstein
Occurs when light strikes the surface of a metal and electrons are ejected.
Practical uses:
Automatic
door openers

20. Photoelectric Effect: Conclusion
Light not only has wave properties but also has particle properties. These massless particles, called photons, are packets of energy.

21. Example 6.2
Using the frequency calculated in the previous example, calculate the energy, in joules, of a photon emitted by an excited sodium atom. Calculate the energy, in kilojoules, of a mole of excited sodium atoms.

22. Bohr’s Hydrogen Atom: A Planetary Model
Niels Bohr: Proposed
planetary model.
Electrons “orbit” the nucleus like planets around the sun.
NOT current model of atom but used to explain some features of atom.

23. Ground State vs. Excited State
ground state: all electrons in lowest possible energy levels
excited state: an electron that has absorbed energy and moved to a higher energy level
This is a temporary state!!

24. Explanation of Line Spectra & Equation
Niels Bohr
Energy of an electron is quantized: can only have specific values.
Energy proportional to energy level.

25. Explanation of Line Spectra
Electron will drop from excited state to ground state and will emit energy as a photon.

26. Explanation of Line Spectra
Type of photon emitted by electron depends on energy difference of energy levels
Elevel = -RH – 1
(nhi) (nlow)2
AND Elevel = h = hc/
(h: Planck’s constant, x J sec/photon)

27. Flaw in Bohr’s Model Does not explain fine structure of line spectra.
Only works well for 1 electron species (H atom).
Does not explain fine structure of line spectra.

28. Wave-Particle Duality
Light has properties of both WAVES and PARTICLES.
most matter has undetectable wavelengths (1000 kg car at 100 km/hr has  = 2.39 x m)
This work led to the development of the electron microscope

29. Quantum Mechanics Quantum mechanics:
atomic structure based on wave-like properties of the electron
Schrödinger: wave equation that describes hydrogen atom

30. Heisenberg Uncertainty Principle
The exact location of an electron cannot be determined (if we try to observe it, we interfere with the particle)
You can know either the location or the velocity but not both
Electrons exist in electron clouds
and not on specific rings or orbits

31. Quantum Numbers
Four quantum numbers are a mathematical way to represent the most probable location of an electron in an atom
analogy...
state = energy level, n
city = sublevel, l
address = orbital, ml
house number = spin, ms

32. Principal Quantum Number: n
Always a positive integer (1,2, 3…7)
Indicates size of orbital, or how far electron is from nucleus
Similar to Bohr’s energy levels or shells
Larger n value = larger orbital or distance from nucleus

33. The Periodic Table and Shells
n = row number on periodic table for a given element
n = 1
n = 2
n = 3
n = 4
n = 5
n = 6
n = 7

34. Angular Momentum Quantum Number: l
positive integer from zero to n-1
Sublevel within an energy level; indicates shape of orbital
0 = s
1 = p
2 = d
3 = f

35. Types of Sublevels s p d

36. Magnetic Quantum Numbers: ml
integer from -l to +l
Indicates orientation of orbital in space
Orbital = electron containing area

37. Spin Quantum Number: ms
Two values only: + ½ or -½
2 electrons max. allowed in each orbital
(Pauli Exclusion Principle)
Indicates spin of electron; spins of each electron must be opposite

38. Every Electron has four!
REVIEW:
QUANTUM NUMBERS
Every Electron has four!
n ---> level 1, 2, 3, 4, ...
l ---> sublevel 0, 1, 2, ... n - 1
ml ---> orbital -l l
ms ---> electron spin +½ and -½

39. Orbitals No more than 2 e- assigned to an orbital
Orbitals grouped in s, p, d (and f) subshells
s orbitals
p orbitals
d orbitals

40. Capacities of levels, sublevels, and orbitals—see packet

41. Example
Example 6.6 Give the n and l values for the following orbitals:
a. 3p
b. 4s

42. Example
Example 6.8 What are the possible ml values for the following orbitals:
a. 3p
b. 4f

43. Shapes of Atomic Orbitals

44. Shapes of Atomic Orbitals
s = spherical
p = peanut
d = dumbbell (clover)
f = flower

45. Multielectron Atoms
In the hydrogen atom the subshells (sublevels) of a principal energy level or shell are at the same energy level.
Previous Equation:
En = –RH /n2

46. Multielectron Atoms
In a multielectron atom, only the orbitals are at the same energy level: the sublevels are at different energy levels!

47. The increasing energy order of sublevels is generally:
s < p < d < f

48. Overlapping subshells
At higher energy levels, sublevels overlap.
Note:
4s vs. 3d!

49. Introduction to Electron Configuration
Definition: describes the distribution of electrons among the various orbitals in the atom
EOS
Represents the most probable location of the electron!

50. Electron Configurations
The system of numbers and letters that designates the location of the electrons
3 major methods:
Full electron configurations
Abbreviated/Noble Gas configurations
Orbital diagram configurations

51. Full or Complete Electron Configuration (uses spdf)
Uses numbers to designate a principal energy level and the letters to identify a sublevel; a superscript number indicates the number of electrons in a designated sublevel.
EOS

52. Rules for Electron Configurations
The Aufbau principle:
Electrons fill from the lowest energy level to the highest (they don’t skip around)
1s22s22p63s23p64s23d10etc.

53. Pauli Exclusion Principle
No two electrons in the same atom can have the same set of 4 quantum numbers.
That is, each electron has a unique “address”
In other words, the maximum # of electrons an orbital can hold is 2 e- (one with ms = +1/2 and one with ms = -1/2)

54. HUND’S RULE
Orbitals of equal energy in a sublevel must all have 1 electron before the electrons start pairing up
a.k.a “creepy person on the bus rule”
*** also electrons in half-filled orbitals have same spin

55. Why are these incorrect?

56. Why are these incorrect?

57. Why are these incorrect?

58. Full Electron Configuration
Example Notation:
1s2 2s1 (Pronounced “one-s-two, two-s-one”)
A. What does the coefficient mean?
Principle energy level
B. What does the letter mean?
Type of orbital (sublevel)
C. What does the exponent mean?
# of electrons in that orbital

59. Steps to Writing Full Electron Configurations
1. Determine the total number of electrons the atom has (for neutral atoms it is equal to the atomic number for the element).
Example: F
atomic # = # of p+ = # of e- =
2. Fill orbitals in order of increasing energy (see Aufbau Chart).
3. Make sure the total number of electrons in the electron configuration equals the atomic number.

60. Aufbau Chart (Order of Energy Levels)
When writing electron configurations:
d sublevels are n – 1 from the row they appear in
f sublevels are n – 2 from the row they appear in

61. Writing Electron Configurations
Nitrogen:
Helium:
Phosphorous:
Rhodium:
Bromine:
Cerium:

62. Abbreviated/Noble Gas Configuration
i. Where are the noble gases on the periodic table?
ii. Why are the noble gases special?
iii. How can we use noble gases to shorten regular electron configurations?

63. Abbreviated/Noble Gas Configuration
Example: Barium
1. Look at the periodic table and find the noble gas in the row above where the element is.
2. Start the configuration with the symbol for that noble gas in brackets, followed by the rest of the electron configuration.

64. Abbreviated/Noble Gas Configuration
Practice! Write Noble Gas Configurations for the following elements:
Rubidium:
Bismuth:
Arsenic:
Zirconium:

65. Writing Electron Configurations
Another way of writing configurations is called an orbital diagram.
(also called orbital notation or orbital diagrams)
One electron has n = 1, l = 0, ml = 0, ms = + ½
Other electron has n = 1, l = 0, ml = 0, ms = - ½

66. Orbital Diagrams = orbital sublevels
Orbital diagrams use boxes (sometimes circles) to represent energy levels and orbitals. Arrows are used to represent the electrons.
= orbital
sublevels

67. Orbital Diagrams
Don’t forget - orbitals have a capacity of two electrons!! Two electrons in the same orbital must have opposite spin so draw the arrows pointing in opposite directions.
Example: oxygen 1s22s22p4 2p
Increasing Energy  2s 1s

68. Drawing Orbital Diagrams
First, determine the electron configuration for the element.
Next draw boxes for each of the orbitals present in the electron configuration.
Boxes should be drawn in order of increasing energy (see the Aufbau chart).
Arrows are drawn in the boxes starting from the lowest energy sublevel and working up. This is known as the Aufbau principle.
Add electrons one at a time to each orbital in a sublevel before pairing them up (Hund’s rule)
The first arrow in an orbital should point up; the second arrow should point down (Pauli exclusion principle)
Double check your work to make sure the number of arrows in your diagram is equal to the total number of electrons in the atom.
# of electrons = atomic number for an atom

69. Electron Configurations for Nitrogen

70. Electron Configurations for Nickel

71. Lithium
Group 1A
Atomic number = 3
1s22s1 ---> 3 total electrons

72. Beryllium
Group 2A
Atomic number = 4
1s22s2 ---> 4 total electrons

73. Boron
Group 3A
Atomic number = 5
1s2 2s2 2p1 --->
5 total electrons

74. Carbon Group 4A Atomic number = 6 1s2 2s2 2p2 --->
6 total electrons
Here we see for the first time HUND’S RULE.

75. Nitrogen Group 5A Atomic number = 7 1s2 2s2 2p3 --->
7 total electrons

76. Oxygen Group 6A Atomic number = 8 1s2 2s2 2p4 --->
8 total electrons

77. Fluorine Group 7A Atomic number = 9 1s2 2s2 2p5 --->
9 total electrons

78. Neon
Group 8A
Atomic number = 10
1s2 2s2 2p6 --->
10 total electrons
Note that we have reached the end of the 2nd period, and the 2nd shell is full!

79. Exceptions to the Filling Order Rule (Cr, Cu)—these will not be on test!

80. Valence electrons Importance and definition:
Definition: Electrons in the outermost energy levels; they determine the chemical properties of an element.
Write the noble gas configuration...the valence electrons are the ones beyond the core.
Example: Sulfur

81. Valence Electrons and Core Configuration (Shorthand)
What is the shorthand notation for S?
EOS
Sulfur has six valence electrons

82. Configurations of Ions
Cations: Formed when metals lose e– in highest principal energy level.
Example:
(Z = 11) Na
EOS
(Z = 11) Na+

83. Configurations of Ions
Anions: Formed when non-metals gain e– to complete the p sublevel.
Example:
EOS -
Z= 18 Cl-

84. Transition Metals
Transition metals (and p block metals) lose e– from the highest principal energy level (n) FIRST, then lose their d electrons!
Zr: [Kr] 5s24d2
Zr+2 : [Kr] 4d2
EOS

85. Isoelectronic Species
Definition: Ions or atoms that have the same number of electrons
Example: Neon, O2-, F-, Na+, Mg2+, Al3+
all have the same configuration (1s22s22p6) and are isoelectronic

86. Electron Spin and Magnetism
Diamagnetic: NOT attracted to a magnetic field
Paramagnetic: substance is attracted to a magnetic field.
Substances with unpaired electrons are paramagnetic.

87. Examples Mg Cl
Write orbital notation: if it has an unpaired e- it is paramagnetic.

88. Periodic Properties & Trends
Electronegativity
Ability of an atom to pull e- towards itself
Increases going up and to the right
Across a period  more protons in nucleus = more positive charge to pull electrons closer
Down a group  more electrons to hold onto = element can’t pull e- as closely

89. Periodic Properties & Trends
Electronegativity
Ability of an atom to pull e- towards itself
Across a period  more protons in nucleus = more positive charge to pull electrons closer
Down a group  more electrons to hold onto = protons in nucleus can’t pull e- as closely

90. Atomic Radius Definition:
½ experimental distance between centers of two bonded atoms

91. Atomic Radius Trend in a family: Size increases down a group.
(More principal
energy levels)

92. Atomic Radius Trend in a period:
Size decreases across a period, e- more strongly attracted to nucleus.

93. Atomic Radius Transition metals:
Size stays relatively constant across a period; e- added to inner energy level.

94. LLLL: Lower Left, Larger Atoms
Memory Device
LLLL: Lower Left, Larger Atoms

95. Sizes of Ions Li +
, 78 pm
2e and 3 p
Li,152 pm
3e and 3p
CATIONS are SMALLER than the atoms from which they are formed.
Size decreases due to increasing he electron/proton attraction.

96. Sizes of Ions F -
, 133 pm
10 e and 9 p
F, 71 pm
9e and 9p
ANIONS are LARGER than the atoms from which they are formed.
Size increases due to more electrons in shell.

97. Trends in ion sizes are the same
as atom sizes.
Active Figure 8.15

98. First Ionization Energy
Definition: energy required to remove an electron from an atom in the gas phase.
Mg (g) kJ ---> Mg+ (g) + e-

99. First Ionization Energies
Trend in a group:
Decreases going down a group (e- further away; easier to remove)
Trend in a period:
Increases going across a period (e- held more tightly).
EOS

100. Memory Device
LLLL: Lower Left,
Larger Atoms;
Looser electrons

101. Second Ionization Energy
Definition: energy required to remove 2nd electron from an atom in the gas phase. Takes more energy because e- is removed from increasingly positive ion.
Mg (g) kJ ---> Mg+ (g) + e-
Mg+ (g) kJ ---> Mg2+ (g) + e-

102. Electron Affinity Some 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

103. Trends in Electron Affinity
Trend in a group:
Affinity for e- decreases going down a group
Trend in a series or period:
Affinity for e- increases going across a period

104. Note that the trend for E.A. is the SAME as for I.E.!
Electron Affinity
Note that the trend for E.A. is the SAME as for I.E.!

105. Trends in Metallic Properties
Most metallic means easiest loss of electrons!
Metals are on left, nonmetals on right of p.t.

106. A Summary of Periodic Trends
Remember LLLL!!