Atomic Structure and Periodicity

1
Atomic Structure and Periodicity

• Chapter 7

2
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.

3
The electromagnetic spectrum
4
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
5
Show Me Problem

• A red light emits light of about 650 nm
  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
6
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 (hn)
7
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.
8
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.
9
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.
10
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.

11
More on the atomic spectrum of hydrogen
12
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)
Dr Quantum Video
13
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 1.634 x 10-18 Joules

Endo or Exo? Does this make sense?
14
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.

15
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
16
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
17
Section 7.6 Quantum Numbers

• Quantum numbers describe the properties of
  orbitals.

symbol name Values meaning
n Principal quantum number 1, 2, etc Energy level
l Angular momentum 0 to (n-1) Orbital shape
ml Magnetic quantum number l to -l Orientationof orbital
18
n l orbital ml of orbitals
1 0 1s 0 1
2 0 2s 0 1
2 1 2p -1, 0, 1 3
3 0 3s 0 1
3 1 3p -1, 0, 1 3
3 2 3d -2, -1, 0, 1, 2 5
4 0 4s 0 1
4 1 4p -1, 0, 1 3
4 2 4d -2, -1, 0, 1, 2 5
4 3 4f -3, -2, -1, 0, 1, 2, 3 7
19
Common Orbital Shapes
20
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

21
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.
22
Electron Configurations

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

23
3 Ways for Electron Configurations
Electron Configuration Diagrams
Long-Hand Configurations
Nobel Gas Configurations
24
Electron Configuration Rule Summary

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

25
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
26
Copper and Chromium

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

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

27
Molybdenum and Silver

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

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

28
Europium and Americium

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

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

29
Electron Configuration Diagrams
30
Noble Gas Notation

• Expanded Titanium 1s22s22p63s23p64s23d2
• Noble Gas Titanium Ar 4s23d2
• Use the noble gas that comes before the element
  as a benchmark, then tack on the extra occupied
  orbitals.

31
Periodic Trends Electron Configurations

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

32
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

33
Periodic Trends Atomic Size
Increasing Atomic Size
34
Periodic Trends Ionization Energy
Increasing 1st Ionization Energy
35
Periodic Trends Electron Affinity
Increasing electron affinity
36
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.