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Electromagnetic Spectrum radiation with both electric and magnetic components

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Title: CHEMISTRY The Molecular Science Author: Ellis CBU Last modified by: Phil Silverman, LT MSC USNR Created Date: 5/8/2001 1:38:08 AM Document presentation format – PowerPoint PPT presentation

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Title: Electromagnetic Spectrum radiation with both electric and magnetic components


1
Electromagnetic Spectrumradiation with both
electric and magnetic components
2
Electromagnetic Radiation
  • Electromagnetic wave
  • wavelength
  • frequency
  • amplitude

3
Wavelength Frequency Amplitude
4
Plancks Quantum Theory
  • E hn hc/l
  • where nl c
  • E ? energy
  • h ? Planck's constant
  • n ? frequency
  • c ? speed of light
  • l ? wavelength

5
Photoelectric Effect
  • the emission of electrons by substances, such as
    metals, when light falls on their surfaces.

6
Photoelectric Effect
7
a quantum, unit, of electromagnetic energyLine
Spectrum
Photons
  • spectrum emitted by a luminous gas or vapor
  • appears as distinct lines characteristic of the
    various elements present

8
Emission Spectrum
  • spectrum of bright lines, bands, or continuous
    radiation
  • characteristic of a specific emitting substance

9
Line Emission Spectrum
10
Energy Transitions
  • Ground State - lowest energy state for electrons
    in an atom
  • Excited State - any energy state above the ground
    state

11
Absorption Spectrum
  • light shinning on a sample causes electrons to be
    excited from the ground state to an excited state
  • wavelengths of that energy are removed from
    transmitted spectra

12
Bohr Model for the Hydrogen Atom
  • E -B/n2
  • E (-2.179 ? 10-18 J/part.)
  • (6.022 ? 1023 part./mole)
  • (1 kJ/103 J)/n2
  • (-1312 kJ/mol)(1/n2)
  • where n ? quantum number
  • 1, 2, 3, 4, 5, 6, 7, etc

13
Bohr Model
  • for hydrogen
  • ground state n 1
  • excited state n gt 1

14
Hydrogen Spectrum
  • E (-2.179 ? 10-18 J)(1/ni2 - 1/nf2)
  • where ni initial state quantum number
  • nf final state quantum number

15
Hydrogen Spectrum
16
Line Spectra
  • Lyman series ? ultraviolet
  • n gt 1 ? n 1
  • Balmer series ? visible light
  • n gt 2 ? n 2
  • Paschen series ? infrared
  • n gt 3 ? n 3

17
Quantum Mechanics
  • Heissenberg Uncertainty Principle - it is
    impossible to know precisely both the location
    and momentum of an electron

18
Orbitals
  • region of probability of finding an electron
    around the nucleus
  • 4 types ? s p d f
  • maximum of 2 electrons per orbital

19
Pure Atomic Orbitals
  • shape of orbitals / energy level
  • s spherical 1
  • p dumbbell 3
  • d complex 5
  • f very complex 7

20
Atomic Orbitals
21
Quantum Numbers
22
Quantum Numbers
  • n ? energy level
  • (1, 2, 3, 4, 5, 6, 7, etc.)
  • l ? type of orbital
  • (s l0 p l1 d l2 f l3)
  • ml ? which orbital
  • (one value per orbital -l ? -l1 ? 0 ? l-1 ? l)
  • s ? direction of spin
  • (-1/2 or 1/2)

23
Electrons
24
Electronic Configurationand the Periodic Table
25
Electronic Configurations
  • The shorthand representation of the occupancy of
    the energy levels (shells and subshells) of an
    atom by electrons.

26
Electronic Configuration
  • H atom 1 electron 1s1
  • He atom 2 electrons 1s2
  • Li atom 3 electrons 1s2, 2s1
  • Cl atom 17 electrons
  • 1s2, 2s2, 2p6, 3s2, 3p5

27
Electronic Configuration
  • As atom
  • 33 electons
  • 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p3
  • or
  • Ar 4s2, 3d10, 4p3

28
Electronic Configuration
  • negative ions
  • add electron(s), 1 electron for each negative
    charge
  • S-2 ion
  • (16 2)electrons
  • 1s2, 2s2, 2p6, 3s2, 3p6

29
Electronic Configuration
  • positive ions
  • remove electron(s), 1 electron for each positive
    charge
  • Mg2 ion
  • (12-2)electrons
  • 1s2, 2s2, 2p6

30
Lewis Electron Dot Structures
31
Trends in thePeriodic Table
  • atomic radius
  • ionic radius
  • ionization energy
  • electron affinity

32
Atomic Radius
  • decrease left to right across a period
  • as nuclear charge increases, number of electrons
    increase however, the nucleus acts as a unit
    charge while the electrons act independently,
    pulling electrons towards the nucleus, decreasing
    size

33
Atomic Radius
  • increase top to bottom down a group
  • each additional electron shell shields the
    outer electrons from the nuclear charge
  • Zeff Z - S
  • where Zeff ? effective nuclear charge
  • Z ? nuclear charge, atomic number
  • S ? shielding constant

34
Atomic Radius
  • increases from upper right corner to the lower
    left corner

35
Atomic Radius ofMain Group Elements
36
Atomic Radius ofs-, p-, and d- Block Elements
37
Nuclear Magnetic Resonance
http//mrsec.wisc.edu/edetc/NMR/index.html
The most common nuclei studied using NMR are
H-1, C-13, F-19 and P-31. Most MRI studies
involve the H-1 nuclei in water. When these
nuclei are exposed to an external magnetic
field, their spins can be parallel or
anti-parallel to this external field.
38
Ionic Radius
  • same trends as for atomic radius
  • positive ions smaller than atom
  • negative ions larger than atom

39
Ionic Radius
  • Isoelectronic Series
  • series of negative ions, noble gas atom, and
    positive ions with the same electronic
    configuration
  • size decreases as positive charge of the
    nucleus increases

40
Isoelectronic Series
41
Sizes of Ions and their Neutral Atoms
42
Ionization Energy
  • energy necessary to remove an electron to form a
    positive ion
  • low value for metals, electrons easily removed
  • high value for non-metals, electrons difficult to
    remove
  • increases from lower left corner of periodic
    table to the upper right corner

43
Ionization Energies
  • first ionization energy
  • energy to remove first electron from an atom
  • second ionization energy
  • energy to remove second electron from a 1 ion
  • etc.

44
First Ionization Energies
45
Electron Affinity
  • energy released when an electron is added to an
    atom
  • same trends as ionization energy, increases from
    lower left corner to the upper right corner
  • metals have low EA
  • nonmetals have high EA

46
Electron Affinities
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