1. Honors Chemistry Chapter 5
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Max Planck (Early 1900’s)
Studied Radiation emitted by solid bodies heated to incandescence.
Thought of the day: Matter could absorb or emit any quantity of energy.
He could not explain his results based on this!!!
So, he proposed that energy can be gained or lost only in whole number intervals of h.
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Planck’s Constant
Transfer of energy is quantized, and can only occur in discrete units, called quanta.
E = change in energy, in J
h = Planck’s constant,  1034 J s
 = frequency, in s1
 = wavelength, in m
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Energy is Quantized
Discrete units of h.
Each small “packet” of energy is called a Quantum.
Energy is transferred in Whole Quanta.
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Surprise!
Energy seems to have particle-like
properties!
Before - energy always assumed to be
continuous.
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Albert Einstein
Proposed that electromagnetic Radiation is itself Quantized, that is,
It can be viewed as a stream of Particles called Photons.
Energy of a photon:
E = h = h c/
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Energy and Mass
Also, Einstein proposed that
Energy has mass
E = mc2
E = energy
m = mass
c = speed of light
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Dual Nature of Light
Electromagnetic radiation exhibits:
1) Wave Properties
2) Particulate Properties
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9. Figure 7.4 Dual Nature of Light
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Light
thought to be purely wavelike
was found to have particulate properties.
Matter
thought to be purely particulate
Does it exhibit wave properties?
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Wavelength and Mass
de Broglie’s Equation (1923): Allows one to calculate an apparent wavelength for a particle.
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Continuous spectrum: Contains all the wavelengths of light.
Line (discrete) spectrum: Contains only some of the wavelengths of light.
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13. Figure 7.6 A Continuous Spectrum (a) and A Hydrogen Line Spectrum (b)
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Sample of H2 gas (H—H)
Introduce a high energy spark
H2 molecules absorb energy
Some of the H—H bonds break
Resulting H atoms are “EXCITED”,
i.e. contain excess energy.
They will eventually “relax” & will release excess energy by emitting light of various wavelengths.
 LINE SPECTRUM
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Line Spectrum
Line spectrum results because only certain energies are allowed for the electron in H atom.
That is, energy of electron in H atom is QUANTIZED.
E = h ν = h c/ λ
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16. Figure 7.7 A Change between Two Discrete Energy Levels
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If any energy were allowed then we would see a Continuous Spectrum (a) and When only certain energies are possible we see only a discrete Line Spectrum (b)
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Niels Bohr (1913)
Developed Quantum Model for the Hydrogen Atom.
The Electron in a Hydrogen Atom moves around the nucleus only in certain allowed circular orbits.
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Figure 7.8 Electronic Transitions in the Bohr Model for the Hydrogen Atom
♦ He calculated the radii for the
allowed circular orbits.
♦ Only certain electron energies
allowed .
♦ Energy levels consistent with the
Hydrogen line-emission spectrum.
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The Bohr Model
Ground State: The lowest possible energy state for an atom (n = 1).
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TWO IMPORTANT POINTS
Bohr Model correctly fits Quantized Energy Levels of the H-atom.
Postulates only certain allowed circular orbits.
As electron is brought closer to the nucleus, Energy is released from the system.
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Bohr’s Model
Appeared promising.
Calculations worked well for hydrogen.
Didn’t work when applied to other atoms.
Something fundamentally incorrect.
Important for its introduction of the concept of Quantization of energy in atoms.
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23. Quantum Mechanical Model of the Atom
Totally different approach was needed.
Three physicists: Heisenberg, de Broglie, & Schrodinger.
Emphasizes the wave properties of an electron.
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24. For the Electron in a Hydrogen Atom
Electron bound to the nucleus.
Similar situation of only certain allowable “Electron Waves.”
Modeled by Schrödinger
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25. Solution of Schrodinger Equation for the Hydrogen Atom
Atomic Orbitals:
space that encloses 90% of the total electron probability.
Wave function for an electron in the Hydrogen atom.
Each electron described by 4 different quantum numbers.
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26. Section 7.6 Quantum Numbers (QN)
Four different quantum numbers.
Three (n, l, ml) specify the wave function that gives the probability of finding the electron at various pts. in space.
Fourth (ms) specifies the electron”spin”.
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Quantum Numbers (QN)
1. Principal QN (n = 1, 2, 3, . . .)
- related to size and energy of the orbital.
- Shell Number
- larger n, then higher energy
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28. 2. Angular Momentum Quantum Number
(l = 0 to n  1)
- relates to shape of the orbital.
- Subshell
l = 0 s subshell
l = 1 p subshell
l = 2 d subshell
l = 3 f subshell
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29. 3. Magnetic Quantum Number
(ml = l to l)
- relates to orientation of the orbital in space relative to other orbitals.
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SHAPES Two Representations of the Hydrogen 1s, 2s, and 3s Orbitals
Node – Area of Zero Probability.
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Figure 7.14 The Boundary Surface Representations of All Three 2p Orbitals
No 1p orbital.
In 2p, two lobes separated by a node at the nucleus.
Labeled according to orientation.
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32. Figure 7.16 The Boundary Surfaces of All of the 3d Orbitals
No 1d or 2d orbitals.
Five 3d orbital
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Figure 7.17 Representation of the 4f Orbitals in Terms of Their Boundary Surfaces
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34. 4. Electron Spin Quantum Number
(ms = +1/2, 1/2)
- relates to the spin states of the electrons.
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35. Electron Spin & Pauli Exclusion Principle
In a given atom, no two electrons can have the same set of four quantum numbers (n, l, ml, ms).
Therefore, an orbital can hold only two electrons, and they must have opposite spins.
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Aufbau Principle
As protons are added one by one to the nucleus to build up the elements, electrons are similarly added to these hydrogen-like orbitals.
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Hund’s Rule
The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli principle in a particular set of degenerate orbitals.
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Greatest triumph
of quantum mechanical model
Is its ABILITY
to account for the
arrangement of elements in the
Periodic Table.
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39. Write the Electronic Configuration for the first 18 elements.
--- full electronic configuration and
--- the noble gas configuration and
--- the orbital diagram.
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Figure 7.26 The Orbitals Being Filled for Elements in Various Parts of the Periodic Table
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Figure 7.24 The Electron Configurations in the Type of Orbital Occupied Last for the First 18 Elements
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42. Figure 7.25 Electron Configurations for Potassium Through Krypton
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Know exceptions of Cu & Cr
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44. 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)
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