Unit 3: The Quantum Model

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Unit 3 The Quantum Model

• Chapter 4
• Starting at back of chapter, then well come
  back to the easier content

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Electrons as Waves Particles

• Where do electrons exist?
• How can you find the exact location and position
  of one, individual electron?
• Electrons are detected by their interaction with
  photons
• They have the same energy, so any attempt to
  locate an e- with a photon knocks the e- off its
  course.
• Heisenberg Uncertainty Principle impossible to
  determine simultaneously both the position and
  velocity of an electron or any other particle

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Schrodinger Wave Equation

• Schrodinger hypothesized that e-s have both a
  dual wave-particle nature
• Quantum Theory describes mathematically the
  wave properties of electrons and other very small
  particles
• Gives the probability of finding an electron at a
  given place around the nucleus
• Orbitals suggested that e-s exist here instead
  of defined orbits

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• Think of a US map listing all the different zip
  codes and their locations
• - Merely s that refer to the positions of
  different postal zones
• Just like an atom, quantum numbers, depict
  positions, and therefore energy levels of
  different e-s in the atom.
• Notice that no two postal codes are the same,
  neither does an atom have the same set of quantum
  numbers.

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Principle Quantum Number

• Referred to as n
• -has integral values of 1, 2, 3,.
• -as n increases, the orbital gets larger
• Sometimes referred to as shells
• -as n increases, more time is spent away from the
  nucleus.
• -as n increases, the e- has a higher energy

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• N contains a certain of sublevels
• Example
• So if n 2, it contains two sublevels, s and p
  

Value of n 1 2 3 4
Type of sublevels s s,p s, p, d s, p, d, f
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Angular Momentum (Azimuthal) Quantum

• Symbolized by L
• -has integral values from 0 to n 1.
• -defines the shape of the orbital
• -the value of L for each orbital is designated by
  the letters, s, p, d, f, which correspond to
  the
• values of 0, 1, 2, 3

Value of L 0 1 2 3
Letters Used s p d f
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Magnetic Quantum Number

• Symbolized by mL
• Example has integrals values between L and - L,
  including 0.
• describes the orientation of the orbital in space
• Example L d
• there are five different orientations that
  correspond to the values, -2, -1, 0,1, 2

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Spin Quantum Number

• Symbolized ms
• Only two possible values, ½ - ½
• Orbital can hold a maximum of two electrons,
  which must have opposite spins

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Electron Configuration

• This is the arrangement of electrons in an atom
• Rules that must be followed
• Aufbau Principle an e- occupies the lowest
  orbital that can receive it.
• Pauli Exclusion Principle no 2 e-s in the same
  atom can have the same set of quantum s
• Hunds Rule orbitals of equal energy are each
  occupied by one electron before any orbital is
  occupied by a second electron, and they must have
  parallel (same) spins

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Grab your blank periodic tables from your unit
plan and your colored pencils.
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Electron Configuration

• 1. The arrangement of electrons in an atom.
• 2. Electrons like to assume arrangements in their
  ground states, because they want the lowest
  possible energy.
• 3. The electron configuration can be described
  pictorially drawn
• denotes the number of electrons in
  orbital or subshell
• 1s1
• denotes n denotes l
• The orbital diagram that shows the spin of the
  electron is

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• Noble Gas Notation
• Example Mg Ne3s2
• Electron Configuration
• Noble Gas Notation
• Orbital Diagram
• Excited State vs. Ground State
• Ions in Electron configuration, noble gas
  notation, and orbital diagrams
• Cr vs Cu
• Isoelectronic

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Valence electrons Lewis dot structures

• Valence electrons are the outermost s and p
  electrons.
• Can never be more than 8.
• These are the electrons used in bonding!!!

• Lewis dot structures show the distribution of
  valence electrons.
• Also cant be more than 8.

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Properties of Light

• Light a form of radiant energy consisting of
  electromagnetic waves that travel freely through
  space
• Electromagnetic radiation form of energy that
  exhibits wavelike behavior as it travels through
  space
• All forms of electromagnetic radiation form the
  electromagnetic spectrum

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• Visible Light
• Features
• Wavelength, Frequency Energy

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Wavelength Frequency

• Wavelength ? distance between corresponding
  points on adjacent waves
• Frequency ? the number of waves that pass a
  given point in a specific time, usually one
  second

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Relating frequency and wavelength

• Use the equation to relate frequency and
  wavelength
• ? is inversely proportional to ?, so in other
  words as the wavelength of light decreases, its
  frequency increases or vice versa.

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The Photoelectric Effect

• Refers to the emission of electrons from a metal
  when light shines on the metal.

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Light as Particles

• Planck proposed that objects emitted energy in
  small, specific amounts called quantum.
• This is the amount of energy that can be lost or
  gained by an atom
• Planck suggested a relationship between a quantum
  of energy and the frequency of radiation

E h v
where h 6.626 x 10-34 J s
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Light having a dual wave-particle like nature

• Einstein expanded on Plancks theory by
  introducing the concept of light have a dual
  wave-particle like nature
• Each particle of light carries a quantum of
  energy, called photons
• Ephoton hv
• Putting Einstein and Planck together
• E mc2
• E hc/?

Solve For m
mc2 hc/?
mc h/ ?
(1/c) mc2 hc/ ? (1/c)
(1/c) mc h/ ? (1/c)
(1/c) mc2 hc/ ? (1/c)
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Continuous Spectrum vs. Line Spectrum

• Continuous Spectrum when white light is passed
  through a prism and all the wavelengths of
  visible light are seen.
• Line Spectrum when the emission spectrum of a
  certain gas is passed through a prism, only bands
  of certain wavelengths are seen.

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Hydrogen-Atom Line Emission Spectrum

• Passed current through a tube containing hydrogen
  gas
• Narrow beam of light passed through prism and a
  series of frequencies or wavelengths were seen.

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• Scientists figured that since only specific
  frequencies of light were emitted then the energy
  differences between the atoms energy were fixed.
• This is what lead Bohr to believe that a hydrogen
  atom exists only in very specific energy states

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• These are additional lines that were discovered
  in the ultraviolet and infrared regions of
  hydrogens line spectrum

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What Bohr Proposed

• 1. The electron on the hydrogen atom can exist
  only in certain spherical orbits.
• 2. As the distance from the nucleus increases,
  the energy of an electron in that orbit
  increases.
• 3. The closest orbit (energy level) is called the
  ground state. Higher energy levels are called
  excited states.
• 4. When an electron falls from a higher energy
  level to a lower energy level, it emits a
  definite amount of energy that is equal to the
  difference in the energy of the two levels.

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Bohrs Model

• ?Ephoton  
• energy of level nfinal -energy of level ninitial