Electromagnetic Radiation

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Electromagnetic Radiation

• Frequency (?, Greek nu) Number of peaks that
  pass a given point per unit time.
• Wavelength (?, Greek lambda) Distance from one
  wave peak to the next.
• Amplitude Height measured from the center of the
  wave. The square of the amplitude gives intensity.

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Electromagnetic Radiation
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Electromagnetic Radiation

• Speed of a wave is the wavelength (in meters)
  multiplied by its frequency in reciprocal
  seconds.
• Wavelength x Frequency Speed
• ? (m) x ? (s1) c (m/s1)

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Electromagnetic Radiation

• The red light in a laser pointer comes from a
  diode laser that has a wavelength of about 630
  nm. What is the frequency of the light?
• c 2.9979 x 108 m/s1

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• Radiation refers to any form of energy that
  travels in all directions from a single source.
• Radiation is NOT associated with nuclear decay.
  This form of radiation is only a very small
  segment of the spectrum.
• Wave frequencies Frequency of a radio wave.

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Figure 7.4
Different behaviors of waves and particles.
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The diffraction pattern caused by light passing
through two adjacent slits.
Figure 7.5
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Electromagnetic Radiation
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Figure 7.7
The line spectra of several elements.
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Atomic Spectra

• Atomic spectra Result from excited atoms
  emitting light.
• Line spectra Result from electron transitions
  between specific energy levels.

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Atomic Spectra 01
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Atomic Spectra 02

• Blackbody radiation is the visible glow that
  solid objects emit when heated.
• Max Planck (18581947) proposed the energy is
  only emitted in discrete packets called quanta.
• The amount of energy depends on the frequency

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Atomic Spectra 03

• Albert Einstein (18791955)
• Used the idea of quanta to explain the
  photoelectric effect.
• He proposed that light behaves as a stream of
  particles called photons.

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Atomic Spectra 05

• For red light with a wavelength of about 630 nm,
  what is the energy of a single photon and one
  mole of photons?

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WaveParticle Duality 01

• Louis de Broglie (18921987) Suggested waves
  can behave as particles and particles can behave
  as waves. This is called waveparticle duality.

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WaveParticle Duality 02

• How fast must an electron be moving if it has a
  de Broglie wavelength of 550 nm?
• How fast must an electron be moving if it has a
  de Broglie wavelength of 630 nm?
• me 9.109 x 1031 kg

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Quantum Mechanics 01

• Niels Bohr (18851962) Described atom as
  electrons circling around a nucleus and concluded
  that electrons have specific energy levels.
• Worked only for the H atom
• Idea that electrons are found in quantized energy
  agreed with the data for the line spectra of
  other atoms.

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Quantum Mechanics 03

• Werner Heisenberg (19011976) Showed that it is
  impossible to know (or measure) precisely both
  the position and velocity (or the momentum) at
  the same time.
• The simple act of seeing an electron would
  change its energy and therefore its position.

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Quantum Mechanics 04

• If an electron has been located to within 0.10 Å,
  what will be the uncertainty in the velocity?
• 1 Å 100 pm 0.10 nm me 9.11 x 1031 kg

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Quantum Mechanics 05

• Erwin Schrödinger (18871961) Developed a
  compromise which calculates both the energy of an
  electron and the probability of finding an
  electron at any point in the molecule.
• This is accomplished by solving the Schrödinger
  equation, resulting in the wave function, ?.
  

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Quantum Numbers 01

• Wave functions describe the behavior of
  electrons.
• Each wave function contains three variables
  called quantum numbers
• Principal Quantum Number (n)
• Angular-Momentum Quantum Number (l)
• Magnetic Quantum Number (ml)

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Quantum Numbers 02

• Principal Quantum Number (n) Defines the size
  and energy level of the orbital. n 1, 2, 3,
  ???
• As n increases, the electrons get farther from
  the nucleus.
• As n increases, the electrons energy increases.
• Each value of n is generally called a shell.

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Quantum Numbers 03

• Angular-Momentum Quantum Number (l) Defines the
  three-dimensional shape of the orbital.
• For an orbital of principal quantum number n, the
  value of l can have an integer value from 0 to n
  1.
• This gives the subshell notation l 0 s
  orbital l 1 p orbital l 2 d orbital
  l 3 f orbital l 4 g orbital

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Quantum Numbers 04

• Magnetic Quantum Number (ml) Defines the spatial
  orientation of the orbital.
• For orbital of angular-momentum quantum number,
  l, the value of ml has integer values from l to
  l.
• This gives a spatial orientation ofl 0 giving
  ml 0 l 1 giving ml 1, 0, 1l 2 giving
  ml 2, 1, 0, 1, 2, and so on...

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Quantum Numbers 11

• Spin Quantum Number
• The Pauli Exclusion Principle states that no two
  electrons can have the same four quantum numbers.

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Quantum Numbers 05
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Quantum Numbers 06

• Why cant an electron have the following quantum
  numbers?
• (a) n 2, l 2, ml 1 (b) n 3, l 0, ml
  3
• (c) n 5, l 2, ml 1
• Give orbital notations for electrons with the
  following quantum numbers
• (a) n 2, l 1, ml 1 (b) n 4, l 3, ml
  2
• (c) n 3, l 2, ml 1

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Electron Radial Distribution 01
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Electron Radial Distribution 02

• s Orbital Shapes

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Electron Radial Distribution 03

• p Orbital Shapes

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Electron Radial Distribution 04

• d and f Orbital Shapes

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Effective Nuclear Charge 01

• Electron shielding leads to energy differences
  among orbitals within a shell.
• Net nuclear charge felt by an electron is called
  the effective nuclear charge (Zeff).

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Effective Nuclear Charge 02

• Zeff is lower than actual nuclear charge.
• Zeff increases toward nucleus ns gt np gt nd gt
  nf
• This explains certain periodic changes observed.

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Effective Nuclear Charge 03
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Electron Configuration of Atoms 01

• Pauli Exclusion Principle No two electrons in an
  atom can have the same quantum numbers (n, l, ml,
  ms).
• Hunds Rule When filling orbitals in the same
  subshell, maximize the number of parallel spins.

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Electron Configuration of Atoms 01

• Rules of Aufbau Principle
• Lower n orbitals fill first.
• Each orbital holds two electrons each with
  different ms.
• Half-fill degenerate orbitals before
  pairingelectrons.

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Electron Configuration of Atoms 01
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Electron Configuration of Atoms 03

• Li ?? ? 1s2 2s1
• 1s 2s
• Be ?? ?? 1s2 2s2
• 1s 2s
• B ?? ?? ? 1s2 2s2 2p1
• 1s 2s 2px 2py 2pz
• C ?? ?? ? ? 1s2 2s2 2p2
• 1s 2s 2px 2py 2pz

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Electron Configuration of Atoms 04

• N ?? ?? ? ? ? 1s2 2s2 2p3
• 1s 2s 2px 2py 2pz
• O ?? ?? ?? ? ? 1s2 2s2 2p4
• 1s 2s 2px 2py 2pz
• Ne ?? ?? ?? ?? ?? 1s2 2s2 2p5
• 1s 2s 2px 2py 2pz
• S Ne ?? ?? ? ? Ne 3s2 3p4
• 3s 3px 3py 3pz

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Electron Configuration of Atoms 05

• Give the ground-state electron configurations
  for
• Ne (Z 10) Mn (Z 25) Zn (Z 30)
• Eu (Z 63) W (Z 74)
• Identify elements with ground-state
  configurations
• 1s2 2s2 2p4 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 5s2
  4d6
• 1s2 2s2 2p6 Ar 4s2 3d1 Xe 6s2 4f14 5d10 6p5

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Electron Configuration of Atoms 06
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Electron Configuration of Atoms 07
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Electron Configuration of Atoms 08

• Anomalous Electron Configurations Result from
  unusual stability of half-filled full-filled
  subshells.
• Chromium should be Ar 4s2 3d4, but is Ar 4s1
  3d5
• Copper should be Ar 4s2 3d9, but is Ar 4s1
  3d10
• In the second transition series this is even more
  pronounced, with Nb, Mo, Ru, Rh, Pd, and Ag
  having anomalous configurations (Figure 5.20).