Download VIII. Other Types of Notations or Configurations

Unit 10 :
Light, Quantum Theory and Electron Configurations
“Where it all starts to come together”
I. Review of Previous Atomic
Structure Theories
• A. Dalton
– 1. All elements are composed of atoms--indivisible and
indestructible particles
– 2. All atoms of the same element are exactly alike-- they
all have the same mass
– 3. Atoms can physically mix or chemically combine in
simple whole # ratios
– 4. Reactions occur when atoms separate, join, or
rearrange
JJ Thomson
• B. J.J. Thompson
– 1. First to theorize the existence of the electron.
– 2. Atomic Model:
• Discovered the e- from work with cathode rays
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“Plum Pudding Model”
Rutherford
• C. Rutherford
– 1. First to theorize the existence of the positive
nucleus.
– 2. Atomic Model:
• Gold Foil Experiment with “Alpha particles”
• “It is about as incredible as if you had fired a 15-inch shell at
a piece of tissue paper and it came back and hit you.” -ER
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Rutherford
Bohr
• D. Bohr
– 1. First to theorize the existence of orbiting
electrons
– 2. Atomic Model:
• “Planetary Model”
Present Model
• E. Present Model
– 1. Names this model is known by:
• A. Charge Cloud Model
• B. Orbital Model
• C. Quantum Mechanical Model
– 2. Location of electrons
• Outside of nucleus
• “probable locations of electrons”
– Electrons are WAVES!
II. The Dual Nature of Light
• A. Modern Theory of Light
– Light has a dual nature-can behave like waves or particles
• B. Light as Waves
– 1. Characteristics of Waves
• A. Wavelength (  “lambda”)
– Distance between any two points in a wave (m)
• B. Frequency (  “nu”)
– # of peaks that pass a given point each second
(cycles/sec  Hz)
II. The Dual Nature of Light
• C. Wave Velocity (speed)
– Distance that a peak travels in a unit of time (1 second)
Relationship:  = c
c = 3.00*108 m/s
-or-
 = c/
Real-World Application
• Favorite Radio Station?
II. The Dual Nature of Light
• C. Discoveries
– 1. Planck’s Theory
• Predicted how the spectrum of light changes with temp
• Energy absorbed or emitted is quantized in nature (See IV. A)
• E = h
(h = Planck’s Constant = 6.626*10-34 J*s)
– 2. Photoelectric Effect
• Einstein proposed light consists of small packets of energy called
photons
• Light behaves like a particle each photon carries a certain amount
of energy (given by Planck’s equation)
II. The Dual Nature of Light
– 3. Compton
• Discovered and concluded that light has properties of both
particles and waves
– 4. Louis de Broglie
• Stated that particles of matter should behave like waves and
exhibit a wavelength (and vice-versa)
• Most matter has undetectable wavelengths
• Related mass, wavelength, and velocity of matter
– =h/m
• Led to the development of the electron microscope
II. The Dual Nature of Light
– 5. Heisenberg’s Uncertainty Principle
• Position and momentum of a moving object can’t be
simultaneously measured and known exactly
• To locate an electron you must strike it with a photon or another
particle
III. The Emission and Absorption of
Radiation
• A. Continuous Spectrum
– 1. Produced by:
• Passing a band of “white light” through a prism
– 2. Looks like:
• The full array of colors (ROYGBIV)
– Red = longest wavelength
– Violet = shortest wavelength
III. The Emission and Absorption of
Radiation
• B. Bright-line Spectrum
– 1. Produced By:
• Passing light through a prism when elements/compounds are
heated
– 2. Looks Like:
• Distinct wavelengths of light given off called “spectral lines”
• Unique for each element (Element DNA)
– Atomic Line Spectra
The electromagnetic spectrum
• Complete range of wavelengths and frequencies
• Mostly invisible
IV. Quantum Theory or Mechanics
(Wave Mechanics)
• A. What does the word “quantized” mean?
– Divide energy into small, but measurable amounts
“Continuous”
“Quantized”
IV. Quantum Theory or Mechanics
(Wave Mechanics)
• B. How does the concept of “quantized”
relate to energy levels?
n=4
n=3
n=2
n=1
IV. Quantum Theory or Mechanics
(Wave Mechanics)
• C. Areas electrons can exist are divided
into:
– 1. Energy levels
– 2. Subshells
– 3. Orbitals
• D. Orbital—Definition
– Region within a sublevel or an energy level
where an e- can be found (homes for e-)
IV. Quantum Theory or Mechanics
(Wave Mechanics)
• E. Ground vs. Excited States
– Ground state lowest energy an e- occupies
– Excited state when an e- is in a higher energy
state
V. Breakdown of the Electron Cloud:
Hotel Analogy
• A. Floor
– 1. Possibilities
• “Principle quantum number”
• Corresponds to energy level/row on Periodic Table
– 2. Like a ENERGY LEVEL in an atom.
• B. Wing or Corridor
– 1. Possibilities
• Names:
– A. s, p, d, f
– B. Floor each one starts on “s”
– 2. Like SUBLEVELS in an atom.
V. Breakdown of the Electron Cloud:
Hotel Analogy
• C. Room
– 1. Possibilities
• Names:
– A. px, py, pz
– B. How many on each corridor number of orbitals = n2
– 2. Like an ORBITAL in an atom.
• D. People
– 1. How many in a room 2 e– 2. Like a PAIR OF ELECTRONS in an atom.
– 3. How do we indicate the different people?
•

VI. Breakdown of the Electron Cloud:
Energy Levels, Sublevels, and Oribitals
• A. Principle Energy Level (n)
– 1. Describes:
• # of principle energy level (row on P.T.)
– 2. Energy Order:
• 1, 2, 3, 4, …
– 3. Number of electron in each energy level
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•
•
•
A. n = 1 2 eB. n = 3 8 eC. n = 3 18 eD. General Formula  2 n2
VI. Breakdown of the Electron Cloud:
Energy Levels, Sublevels, and Oribitals
• B. Energy Sublevels
– 1. Types of sublevels
Sublevel
s
Shape
# of each Energy Levels?
n = 1, 2, 3, 4
p
n = 2, 3, 4
d
n = 3, 4
f
n=4
VI. Breakdown of the Electron Cloud:
Energy Levels, Sublevels, and Oribitals
– 2. Energy order of sublevels
• n =1 < 2 < 3 < 4
– 3. General formula for number of sublevels at
each energy level
• # sublevels = principle quantum #
– 4. Orientation of sublevels in space
•s
• px, py, pz
•d
VI. Breakdown of the Electron Cloud:
Energy Levels, Sublevels, and Oribitals
• C. Orbitals
– 1. Maximum number of electrons in an orbital
• 2 e-
– 2. General formula for number of electrons in a
principle energy level
• 2 n2  maximum number of e-
– 3. Electron pair—definition
• Only two e- can occupy an orbital (within a sublevel) at a time.
They must have opposite spins.
• D. See Summary Pg. 11
– Complete table on Pg. 12
VII. Arrangement or Configuration of
Electrons
• A. Order is determined by two basic principles:
– Pauli Exclusion Principle (2e-/orbital)
– Hund’s Rule (half-filled before completely filled orbitals)
• B. Energy Order—See Pg. 14
• C. Memory device for energy order—see attachment
VII. Arrangement or Configuration of
Electrons
• D. Regular Electron Configuration
– 1. Example Notation
• 1s2 2s1 (Pronounced “one-s-two, two-s-one”)
• A. What does the coefficient mean?
– Principle Quantum #
• B. What does the letter mean?
– Type of orbital (subshell)
• C. What does the exponent mean?
– # of electrons in that orbital
VII. Arrangement or Configuration of
Electrons
• E. How to write e- configurations
• F. Examples of e- configurations
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–
–
–
1. He
2. N
3. Na
4. Ca
1s2
1s2 2s2 2p3
1s2 2s2 2p6 3s1
1s2 2s2 2p6 3s2 3p6 4s2
• Practice Time! (Pg. 16)
VIII. Other Types of Notations or
Configurations
• A. Orbital Notation or Configuration
– 1. How are each of the following represented?
• A. Orbital
• B. Electron
VIII. Other Types of Notations or
Configurations
– 2. What does each type of sublevel look like in
orbital notation?
• A. s
• B. p
• C. d
• D. f
• *Note* Put brackets around each set of orbitals!
VIII. Other Types of Notations or
Configurations
– 3. How is each sublevel filled?
• A. Hund’s Rule
– Every orbital in a subshell (s, p, d, f) is singly occupied with
one e- before any one orbital is doubly occupied
• B. Pauli Exclusion Principle
– In order for two electrons to occupy the same orbital, they
must have opposite spins
– 4. How to write an orbital configuration (Pg. 19)
VIII. Other Types of Notations or
Configurations
– 5. Examples of Orbital Configurations
• A. He
• B. N
• C. Na
• D. Ca
VIII. Other Types of Notations or
Configurations
• B. Shorthand Configuration
– 1. Where are the noble gases on the periodic
table?
• Group VIII (far right)
– 2. Why are the noble gases special?
• Noble gases have 8 e- in their outer shell
• Very stable (least reactive)
– 3. How can we use noble gases to shorten the
regular e- configurations?
• Ex. Ba
VIII. Other Types of Notations or
Configurations
– 4. How to write a shorthand configuration (Pg. 22)
– 5. Examples of shorthand configuration
• A. Na
• B. Mn
• C. Ca
IX. Relating Electron Configurations to
the Periodic Table
• A. Location of the final electrons in each
element:
–
–
–
–
1. s (Groups I, II)
2. p (Groups IIIA – VIIIA)
3. d (Transition metals)
4. f (Lanthanides, Actinides)
• How to write configurations using the P.T.
X. Valence Electrons and Lewis
(Electron Dot) Structures
• A. Definition
– Valence Electrons
• Electrons that are in the outermost principle energy
level (bonding electrons)
• “Where the action is” in Chemistry
• B. Outermost principle energy level is
called the valence shell
– must include at least 1 electron
X. Valence Electrons and Lewis
(Electron Dot) Structures
• C. Inner electrons are called core or kernel
electrons
• D. Examples
– Be: 1s2 2s2
– O: 1s2 2s2 2p4
– Sc: 1s2 2s2 2p6 3s2 3p6 4s2 3d1
2 valence e4 valence e2 valence e-
X. Valence Electrons and Lewis
(Electron Dot) Structures
• E. How to find valence e- easily?
– The highest energy level (n = ?): # of e- here!
• F. Representing valence e- in atoms using Lewis
or Electron Dot structures
– 1. Chemical Symbol
– 2. Each dot represents valence e– 3. “s” orbital valence e- double up on top, then single egoing clockwise
X. Valence Electrons and Lewis
(Electron Dot) Structures
• G. Examples of Lewis Dot Structures
– 1. Hydrogen
– 2. Carbon
– 3. Phosphorus
End of Unit 10!!