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Atomic Structure and Periodicity

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Title: 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 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
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 n(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)
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 -1.633x10-18 Joules

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.

18
(No Transcript)
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
Periodic Trends Electron Configurations
  • Counting down tells what energy orbital.
  • Counting over tells how many electrons.

30
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

31
Periodic Trends Atomic Size
Increasing Atomic Size
32
Periodic Trends Ionization Energy
Increasing 1st Ionization Energy
33
Periodic Trends Electron Affinity
Increasing electron affinity
34
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.
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