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Electromagnetic Radiation

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Title: Electromagnetic Radiation


1
Electromagnetic Radiation
  • Light is a form of energy.
  • Technically, light is one type of a more general
    form of energy called electromagnetic radiation
    and travels in waves.
  • Every wave has four characteristics that
    determine its properties wave speed, height
    (amplitude), length, and the number of wave peaks
    that pass in a given time.
  • Velocity (c) is the speed of light and equals
    2.997925 x 108 m/s (msec-1) in a vacuum.
  • It is constant! All types of light energy travel
    at the same speed.
  • Amplitude (A) is a measure of the intensity of
    the wave (the height of the wave) or
    brightness.
  • Wavelength (l ) is the distance between crests.
  • It is generally measured in nanometers (1 nm
    10-9 m)
  • frequency (n) represents the number of peaks
    passing a point in one second.
  • It is generally measured in Hertz (Hz) where
  • 1 Hz 1 wave/sec 1 sec-1

2
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3
Particles of Light
  • Scientists in the early 20th century showed that
    electromagnetic radiation was composed of
    particles we call photons, particles of light
    energy.
  • Max Planck and Albert Einstein
  • Each wavelength of light has photons that have a
    different amount of energy

The Electromagnetic Spectrum
  • Light passing through a prism is separated into
    all its colors and is called a continuous
    spectrum.
  • The color of the light is determined by its
    wavelength.

4
  • Wavelength and frequency are related by the speed
    of light c ? ?
  • The speed of light, c, is a constant
    therefore, wavelength and the frequency
    are inversely proportional as the wavelength
    increases, the frequency decreases.
  • Planck related the frequencies of EM radiation to
    the energies of vibrational transitions in matter
    to give the equation E h ?
  • where h Plancks constant 6.626 x 10-34 Js
  • This equation represents the energy of a SINGLE
    photon of the given frequency.
  • Energy and frequency are directly proportional
  • High frequency short wavelength high energy
  • Low frequency long wavelength low energy

5
Types of Electromagnetic Radiation
  • Classified by the Wavelength
  • Radio waves l gt 0.01 m
  • low frequency and energy
  • Microwaves 10-4m lt l lt 10-2m
  • Infrared (IR) 8 x 10-7 lt l lt 10-5m
  • Visible 4 x 10-7 lt l lt 8 x 10-7m
  • ROYGBV
  • Ultraviolet (UV) 10-8 lt l lt 4 x 10-7m
  • X-rays 10-10 lt l lt 10-8m
  • Gamma rays l lt 10-10
  • high frequency and energy

6
Electromagnetic Spectrum
7
Lights Relationship to Matter
  • Atoms can acquire extra energy, but they must
    eventually release it.
  • When atoms emit energy, it always is released in
    the form of light.
  • However, atoms dont emit all colors, only very
    specific wavelengths.
  • In fact, the spectrum of wavelengths can be used
    to identify the element!

8
Emission Spectrum
Spectra
9
The Bohr Model of the Atom
  • The Nuclear Model of the atom does not explain
    how the atom can gain or lose energy.
  • Neils Bohr developed a model of the atom to
    explain how the structure of the atom changes
    when it undergoes energy transitions.
  • Bohrs major idea was that the energy of the atom
    was quantized, and that the amount of energy in
    the atom was related to the electrons position
    in the atom
  • quantized means that the atom could only have
    very specific amounts of energy.

10
The Bohr Model of the AtomElectron Orbits
  • In the Bohr Model, electrons travel in orbits
    around the nucleus
  • more like shells than planet orbits.
  • The farther the electron is from the nucleus the
    more energy it has

Each orbit has a specific amount of energy. The
energy of each orbit is characterized by an
integer - the larger the integer, the more energy
an electron in that orbit has and the farther it
is from the nucleus. The integer, n, is called a
quantum number
11
Energy Transitions
  • When an atom gains energy, an electron leaps from
    a lower energy orbit to one that is further from
    the nucleus.
  • However, during that quantum leap it doesnt
    travel through the space between orbits it just
    disappears from the lower orbit and appears in a
    higher orbit!
  • When the electron leaps from a higher energy
    orbit to one that is closer to the nucleus,
    energy is emitted from the atom as a photon of
    light!

12
Ground and Excited States
  • In the Bohr Model of hydrogen, the lowest amount
    of energy hydrogens one electron can have
    corresponds to being in the n 1 orbit we call
    this its ground state.
  • When the atom gains energy, the electron leaps to
    a higher energy orbit we call this an excited
    state.
  • The atom is less stable in an excited state, and
    so it will release the extra energy to return to
    the ground state
  • either all at once or in several steps.

Every hydrogen atom has identical orbits, so
every hydrogen atom can undergo the same energy
transitions. However, since the distances between
the orbits in an atom are not all the same, no
two leaps in an atom will have the same
energy. The closer the orbits are in energy, the
lower the energy of the photon emitted. A lower
energy photon a longer wavelength. Therefore,
we get an emission spectrum that has a lot of
lines that are unique to hydrogen.
13
The Bohr Model of theHydrogen Spectrum
Success and Failure
The mathematics of the Bohr Model very accurately
predicts the spectrum of hydrogen but fails when
applied to multi-electron atoms. It cannot
account for electron-electron interactions.
Enter the Quantum-Mechanical Model of the atom! ?
14
The Quantum-Mechanical Model of the Atom
  • Erwin Schr√∂dinger applied the mathematics of
    probability and the ideas of quantization to the
    physics equations that describe waves resulting
    in an equation that predicts the probability of
    finding an electron with a particular amount of
    energy at a particular location in the atom.
  • The result is a map of regions in the atom that
    have a particular probability for finding the
    electron.
  • An orbital is a region where we have a very high
    probability of finding the electron when it has a
    particular amount of energy.
  • It is generally set at 90 or 95.

15
Orbits vs. OrbitalsPathways vs. Probability
16
The Quantum-Mechanical ModelQuantum Numbers
  • in Schr√∂dingers Wave Equation, there are 3
    integers, called quantum numbers, that quantize
    the energy
  • the principal quantum number, n, specifies the
    main energy level for the orbital
  • each principal energy shell has one or more
    subshells
  • the number of subshells the principal quantum
    number
  • the quantum number that designates the subshell
    is often given a letter
  • s, p, d, f
  • each kind of sublevel has orbitals with a
    particular shape
  • the shape represents the probability map
  • 90 probability of finding electron in that
    region

17
Shells Subshells
18
How does the 1s Subshell Differ from the 2s
Subshell
Probability Maps Orbital Shapes Orbitals
19
Probability Maps Orbital Shapep Orbitals
d Orbitals
20
Subshells and Orbitals
  • The subshells of a principal shell have slightly
    different energies.
  • The subshells in a shell of H all have the same
    energy, but for multielectron atoms the subshells
    have different energies
  • s lt p lt d lt f
  • Each subshell contains one or more orbitals
  • s subshells have 1 orbital
  • p subshells have 3 orbitals
  • d subshells have 5 orbitals
  • f subshells have 7 orbitals
  • Each energy shell and subshell has a maximum
    number of electrons it can hold
  • s 2, p 6, d 10, f 14

21
  • Each orbital may have a maximum of 2 electrons
  • -Pauli Exclusion Principle
  • Electrons spin on an axis generating their own
    magnetic field.
  • When two electrons are in the same orbital, they
    must have opposite spins so that their magnetic
    fields will cancel.

Orbital Diagrams
22
Order of Subshell Fillingin Ground State
Electron Configurations
Start by drawing a diagram putting each energy
shell on a row and listing the subshells, (s, p,
d, f), for that shell in order of energy,
(left-to-right).
1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d
7s
Next, draw arrows through the diagonals, looping
back to the next diagonal each time.
23
The Quantum Mechanical ModelEnergy Transitions
  • As in the Bohr Model, atoms gain or lose energy
    as the electron leaps between orbitals in
    different energy shells and subshells.
  • The ground state of the electron is the lowest
    energy orbital it can occupy.
  • Higher energy orbitals are excited states.
  • Both the Bohr and Quantum Mechanical models
    predict the spectrum of hydrogen very accurately.
  • Only the Quantum Mechanical model predicts the
    spectra of atoms with more than one electron.

24
Electron Configuration of Atoms in their Ground
State
  • the electron configuration is a listing of the
    subshells in order of filling with the number of
    electrons in that subshell written as a
    superscript
  • Kr 36 electrons 1s22s22p63s23p64s23d104p6
  • a shorthand way of writing an electron
    configuration is to use the symbol of the
    previous noble gas in to represent all the
    inner electrons, then just write the last set
  • Rb 37 electrons 1s22s22p63s23p64s23d104p65s1
    Kr5s1

25
Example Write the Ground State Orbital Diagram
and Electron Configuration of Magnesium.
  • Determine the atomic number of the element from
    the Periodic Table
  • This gives the number of protons and electrons in
    the atom
  • Mg Z 12, so Mg has 12 protons and 12 electrons
  • Draw 9 boxes to represent the first 3 energy
    levels s and p orbitals.
  • Add one electron to each box in a set, then pair
    the electrons before going to the next set until
    you use all the electrons.
  • When pairing the electrons, put the arrows in
    opposite directions.
  • Use the diagram to write the electron
    configuration
  • Write the number of electrons in each set as a
    superscript next to the name of the orbital set
  • 1s22s22p63s2 Ne3s2

26
Valence Electrons
  • The electrons in all the subshells with the
    highest principal energy shell are called the
    valence electrons. These electrons can
    participate in chemical bonding. If an atom is in
    the second row, then all the electrons in the
    second row are valence electrons
  • Electrons in lower energy shells are called core
    electrons. In the electron configuration, the
    core electrons are equivalent to a noble gas
    configuration. These electrons do NOT participate
    in chemical bonding
  • Rb 37 electrons 1s22s22p63s23p64s23d104p65s1
  • the highest principal energy shell of Rb that
    contains electrons is the 5th, therefore Rb has 1
    valence electron and 36 core electrons
  • Kr 36 electrons 1s22s22p63s23p64s23d104p6
  • the highest principal energy shell of Kr that
    contains electrons is the 4th, therefore Kr has 8
    valence electrons and 28 core electrons

27
Electron Configurations andthe Periodic Table
28
Electron Configuration fromthe Periodic Table
8A
1A
1 2 3 4 5 6 7
3A
4A
5A
6A
7A
2A
Ne
P
3s2
3p3
P Ne3s23p3 P has 5 valence electrons
29
The Explanatory Power ofthe Quantum-Mechanical
Model
  • The properties of the elements are largely
    determined by the number of valence electrons
    they contain.
  • Since elements in the same column have the same
    number of valence electrons, they show similar
    properties

30
Stable Electron ConfigurationAnd Ion Charge
  • Metals form cations by losing enough electrons to
    get the same electron configuration as the
    previous noble gas.
  • Nonmetals form anions by gaining enough electrons
    to get the same electron configuration as the
    next noble gas.

31
Trends in Ionization Energy
32
Trends in Atomic Size
33
Metallic Character
Metallic Properties Shiny Tend to form
cations Low ionization energies 3 or fewer
valence electrons High conductivity Malleable
(shapeable)
Nonmetallic Properties Dull - no luster Tend to
form anions High ionization energy 4 or more
valence electrons Low conductivity Brittle
34
Trends in Metallic Character
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