Atomic Structure and Periodicity

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Atomic Structure and Periodicity

• Chapter 7

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Section 7.1 Electromagnetic Radiation

• The electromagnetic spectrum organizes waves by
  wavelength, frequency, and energy
• l wavelength- the length from a point on a wave
  to a corresponding point later in the wave.
• n frequency- the number of times a full wave
  cycle passes by a reference point in one second.

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The electromagnetic spectrum
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Relationships between l, n, and E

• Wavelength and frequency are inversely
  proportional (a long wave will take longer to
  pass by a reference point, thus making its
  frequency lower).
• Energy is directly related to frequency. Higher
  frequency waves have more energy than lower
  frequency waves.
• The speed of a wave is constant in a vacuum. (3.0
  x 108 m/s c, speed of light)

l n c
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Show Me Problem

• A red light emits light of about 650 m
  wavelength. What is the frequency of the red
  light?
• 650 nm ? 6.5 x 10-7m
• (6.5x10-7)(n) (3.0 x 108)
• n 4.61 x 1014 Hz

l n c
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Section 7.2 The Nature Of Matter

• Max Plank
• Quantizes energy
• Matter can absorb energy, but only in whole
  number ratios of the term hn.

Planks constant 6.626 x 10-34 Js
EQUATION DE n(hn)
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Einsteins Contribution

• Albert Einstein
• Quantizes Radiation
• E mc2 Energy has mass and velocity.
  Electromagnetic Radiation must be made up of
  particles called photons.

Duality of Light Electromagnetic radiation has
the capacity to behave both as a wave and a
particle.
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DeBroglies Contribution

• Louis DeBroglie
• Determines the Duality of Matter
• If waves act as particles, do particles act as
  waves?
• Set the Einstein and the Plank equations equal to
  each other.
• mc2 E hc/l
• m h/cl
• l h/mn

Duality of Matter Electrons have the capacity
to behave both as a particle and a wave.
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How can we be certain?

• X Ray Crystallography

Different color and shading patters appear as the
electron waves cause diffraction constructive
and destructive interference with the x-rays that
are exposed to the crystal.
X Ray Diffraction Pattern of 2-terphenyl-4-yl-5-p
henyl thiophene (PPPTP)
X-Ray Diffraction Pattern of Beryl.
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Atomic Spectrum of Hydrogen

• Extensively studied by atomic theorists such as
  Bohr.
• High energy sparks cause hydrogen gas molecules
  (H-H) to break apart suddenly, with some
  electrons in higher energy levels than would be
  expected normally.
• As the electrons fall back to their ground
  states, energy is released.
• Each color in the spectrum relates to an electron
  in a different energy level.
• Planks equation can be used to determine the
  color of the light produced or the energy of the
  electron that is being observed.

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More on the atomic spectrum of hydrogen
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Section 7.4 The Bohr Model

• Based upon the study of the Hydrogen Spectrum,
  Bohr designs paths for electrons to travel while
  orbiting the nucleus. ORBITS
• Each orbit corresponded to a different energy
  level.

energy of e-
energy level
nuclear charge (protons)
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Show Me

• An electron in a hydrogen atom moves from energy
  level one to energy level 2. What is the change
  in energy the electron experiences?
• E1 -2.178 x 10-18 (12/12) -2.178 x 10-18
• E2 -2.178 x 10-18 (12/22) -5.445 x 10-19
• DE E2-E1 -5.445 x 10-19-(-2.178 x 10-18)
• DE -1.633x10-18 Joules

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Section 7.5 The Quantum Mechanical Model of the
Atom
Heisenberg
De Broglie
Scrhödinger

• Determine that if the electron acts as a standing
  wave (a wave that is fixed in place), then there
  are only certain orientations for it to exist
  without causing destructive interference with
  itself.

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Heisenberg Uncertainty Principle

• There is a fundamental limitation to just how
  precisely we can know both the position and
  momentum of a particle.
• The more certain you are of the location of an
  electron, the less certain you can be of its
  momentum
• The more certain you are of the momentum of an
  electron, the less certain you can be of its
  position.

Dx Dmn h/4p
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The Wave Equation

• Schrodingers wave equation is used to define the
  location of electrons as waves.
• A wave function is called an orbital.
• ORBITALS ORBITS
• Wave functions are impossible to visualize. We
  picture the electron density map (aka electron
  probability diagram)

HY EY
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Section 7.6 Quantum Numbers

• Quantum numbers describe the properties of
  orbitals.

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Common Orbital Shapes
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Section 7.7 Orbital Shapes

• Areas of high probability are separate by areas
  of low probability. (NODES)
• Degenerate orbitals have different orientation or
  shape but the same ENERGY.
• Lowest available energy level for an electron
  ground state
• Higher energy levels than expected excited
  states

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Section 7.8 Electron Spin and Pauli Principle

• Electrons exhibit a fourth quantum number.
• Electron spin quantum number (ms)
• Values of ½ and -½. Indicates magnetic moment of
  electron. Electrons can only spin in one of 2
  opposite directions.

Pauli exclusion Principle In a given atom, no
two electrons Can have the same set of Quantum
numbers.
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Electron Configurations

• Diagonal Rule
• 5s 5p 5d 5f
• 4s 4p 4d 4f
• 3s 3p 3d
• 2s 2p
• 1s

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3 Ways for Electron Configurations
Electron Configuration Diagrams
Long-Hand Configurations
Nobel Gas Configurations
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Electron Configuration Rule Summary

• Electrons enter lowest energy orbitals first.
• Only two electrons per degenerate orbital.
• Electrons spread out among degenerate orbitals
  before pairing.

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Exceptions to the Configuration Rules

• A fully filled orbital is more stable than a
  partially filled orbital.
• half-filled orbital is more stable than a
  more/less partially filled orbital.

Mo
Ag
Eu
Am
Cr
Cu
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Copper and Chromium

• Cu expected configuration
• 1s22s22p63s23p64s23d9
• Cu actual configuration
• 1s22s22p63s23p64s13d10

• Cr expected configuration
• 1s22s22p63s23p64s23d4
• Cr actual configuration
• 1s22s22p63s23p64s13d5

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Molybdenum and Silver

• Mo expected configuration
• 1s22s22p63s23p64s23d104p65s24d4
• Mo actual configuration
• 1s22s22p63s23p64s23d104p65s14d5

• Ag expected configuration
• 1s22s22p63s23p64s23d104p65s24d9
• Ag actual configuration
• 1s22s22p63s23p64s23d104p65s14d10

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Europium and Americium

• Eu expected configuration
• 1s22s22p63s23p64s23d104p65s24d105p66s24f6
• Eu actual configuration
• 1s22s22p63s23p64s23d104p65s24d105p66s14f7

• Am expected configuration
• 1s22s22p63s23p64s23d104p65s24d105p66s24f145d107s25
  f6
• Am actual configuration
• 1s22s22p63s23p64s23d104p65s24d105p66s24f145d107s15
  f7

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Periodic Trends Electron Configurations

• Counting down tells what energy orbital.
• Counting over tells how many electrons.

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Periodic Trends Activity

• http//academic.pgcc.edu/ssinex/excelets/PT_inter
  active.xls
• Go to this excel sheet and click on the bottom
  tab labeled atom properties

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Periodic Trends Atomic Size
Increasing Atomic Size
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Periodic Trends Ionization Energy
Increasing 1st Ionization Energy
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Periodic Trends Electron Affinity
Increasing electron affinity
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Ion Size

• Negative ions indicate a gain of electrons.
• They are larger than the atom from whence they
  are formed.
• Positive ions indicate a loss of electrons.
• They are smaller than the atom from whence they
  are formed.