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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
E = change in energy, in J
h = Planck’s constant, 6.626 1034 J s
= frequency, in s1
= wavelength, in m
E hhc
= =
Transfer of energy is quantized, and can Transfer of energy is quantized, and can only occur in discrete units, called only occur in discrete units, called quantaquanta..
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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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Figure 7.4
Dual Nature of Light
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Light thought to be purely wavelikewas found to have particulate properties.
Matter
thought to be purely particulateDoes it exhibit wave properties?
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Wavelength and Massde Broglie’s Equation (1923): Allows one to de Broglie’s Equation (1923): Allows one to calculate an apparent wavelength for a calculate an apparent wavelength for a particle.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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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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Figure 7.7A 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 possiblewe 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 theallowed circular orbits.
♦ Only certain electron energiesallowed .
♦ 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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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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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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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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Section 7.6Quantum 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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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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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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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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4. Electron Spin Quantum Number
(ms = +1/2, 1/2)
- relates to the spin states of the electrons.
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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 thearrangement of elements in thePeriodic Table.
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Write the Electronic Configuration for the first 18 elements.
• Write the
--- full electronic configuration and
--- the noble gas configuration and
--- the orbital diagram.
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Figure 7.26The 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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Figure 7.25
Electron Configurations for Potassium Through Krypton
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• Know exceptions of Cu & Cr
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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)