Chapter 13 Electrons in Atoms by HC111123011234

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									   Chapter 13
Electrons in Atoms
       Section 13.1
     Models of the Atom
 OBJECTIVES:
 Summarize the development of
  atomic theory.
       Section 13.1
     Models of the Atom
 OBJECTIVES:
 Explain the significance of
  quantized energies of electrons
  as they relate to the quantum
  mechanical model of the atom.
              Greek Idea
 Democritus
 Matter is made up
  of solid indivisible
  particles
 John Dalton - one
  type of atom for
  each element
     J. J. Thomson’s Model
 Discovered electrons
 Atoms were made of
  positive stuff
 Negative electron
  floating around
 “Plum-Pudding”
  model
    Ernest Rutherford’s Model
 Discovered dense
  positive piece at
  the center of the
  atom- nucleus
 Electrons would
  surround it
 Mostly empty
  space
 “Nuclear model”
        Niels Bohr’s Model
 He  had a question: Why don’t the
  electrons fall into the nucleus?
 Move like planets around the sun.
 In circular orbits at different levels.
 Amounts of energy separate one
  level from another.
 “Planetary model”
    Bohr’s planetary model
 Energy level of an electron
 analogous to the rungs of a ladder
 electron cannot exist between energy
  levels, just like you can’t stand
  between rungs on ladder
 Quantum of energy required to move
  to the next highest level
    The Quantum Mechanical
            Model
 Energy is quantized. It comes in chunks.
 A quanta is the amount of energy needed to
  move from one energy level to another.
 Since the energy of an atom is never “in
  between” there must be a quantum leap in
  energy.
 Erwin Schrodinger derived an equation that
  described the energy and position of the
  electrons in an atom
    The Quantum Mechanical
            Model
 Things that are very small
  behave differently from things
  big enough to see.
 The quantum mechanical
  model is a mathematical
  solution
 It is not like anything you can
  see.
    The Quantum Mechanical
            Model
 Has energy levels for
  electrons.
 Orbits are not circular.
 It can only tell us the
  probability of finding
      an electron a certain distance from
  the nucleus.
    The Quantum Mechanical
            Model
 The atom is found
  inside a blurry
  “electron cloud”
 A area where there is
  a chance of finding
  an electron.
 Draw a line at 90 %
 Think of fan blades
          Atomic Orbitals
 Principal Quantum Number (n) = the
  energy level of the electron.
 Within each energy level, the complex
  math of Schrodinger’s equation
  describes several shapes.
 These are called atomic orbitals -
  regions where there is a high probability
  of finding an electron.
 Sublevels- like theater seats arranged in
  sections
Two
representations
of the hydrogen
1s, 2s, and 3s
orbitals.
Representation of the 2p orbitals. (a) The
 electron probability distributed for a 2p
    orbital. (b) The boundary surface
 representations of all three 2p orbitals.
Representation of the 3d orbitals.
Representation of the 4f orbitals in
 terms of their boundary surface.
            Summary
    # of   Max         Starts at
    shapes electrons   energy level
s    1        2            1

p    3        6            2

d    5        10           3

f    7        14           4
         By Energy Level
 First Energy Level    Second Energy
 only s orbital         Level
 only 2 electrons      s and p orbitals

 1s
     2                   are available
                        2 in s, 6 in p
                            2 6
                        2s 2p
                        8 total electrons
          By Energy Level
 Third energy level     Fourth energy
 s, p, and d             level
  orbitals               s,p,d, and f
 2 in s, 6 in p, and     orbitals
  10 in d                2 in s, 6 in p, 10
     2 6 10
 3s 3p 3d                in d, ahd 14 in f
                         4s
                             24p64d104f14
 18 total electrons
                         32 total electrons
           By Energy Level
 Any more than            The orbitals do
  the fourth and not        not fill up in a
  all the orbitals will     neat order.
  fill up.                 The energy levels
 You simply run            overlap
  out of electrons         Lowest energy fill
                            first.
         Section 13.2
Electron Arrangement in Atoms

 OBJECTIVES:
  Apply the aufbau principle, the
   Pauli exclusion principle, and
   Hund’s rule in writing the electron
   configurations of elements.
         Section 13.2
Electron Arrangement in Atoms

 OBJECTIVES:
  Explain why the electron
   configurations for some elements
   differ from those assigned using
   the aufbau principle.
                         7p         6d
                    7s   6p                      5f
                                    5d
                    6s   5p                      4f
                                    4d
                    5s
Increasing energy



                         4p
                    4s              3d
                         3p
                    3s
                         2p
                    2s

                         Aufbau diagram - page 367
                    1s
      Electron Configurations
 The way electrons are arranged in
  atoms.
 Aufbau principle- electrons enter the
  lowest energy first.
 This causes difficulties because of the
  overlap of orbitals of different energies.
 Pauli Exclusion Principle- at most 2
  electrons per orbital - different spins
     Electron Configuration
 Hund’s  Rule- When electrons
  occupy orbitals of equal energy
  they don’t pair up until they have to.
 Let’s determine the electron
  configuration for Phosphorus
 Need to account for 15 electrons
                         7p          6d
                    7s   6p                        5f
                                     5d
                    6s   5p                        4f
                                     4d
                    5s
Increasing energy



                         4p
                    4s               3d
                         3p   The first two electrons
                              
                    3s
                              go into the 1s orbital
                         2p
                    2s       Notice the opposite
                              spins
                             only 13 more to go...
                    1s
                         7p           6d
                    7s   6p                          5f
                                      5d
                    6s   5p                          4f
                                      4d
                    5s
Increasing energy



                         4p
                    4s                3d
                         3p
                    3s
                         2p    The next electrons
                    2s          go into the 2s orbital
                               only 11 more...

                    1s
                         7p          6d
                    7s   6p                        5f
                                     5d
                    6s   5p                        4f
                                     4d
                    5s
Increasing energy



                         4p
                    4s               3d
                         3p
                    3s
                         2p   • The next electrons go
                    2s          into the 2p orbital
                              • only 5 more...
                    1s
                         7p          6d
                    7s   6p                         5f
                                     5d
                    6s   5p                         4f
                                     4d
                    5s
Increasing energy



                         4p
                    4s               3d
                         3p
                    3s
                         2p   • The next electrons go
                    2s          into the 3s orbital
                              • only 3 more...
                    1s
                         7p              6d
                    7s   6p                              5f
                                          5d
                    6s   5p                              4f
                                          4d
                    5s
Increasing energy



                         4p
                    4s                    3d
                         3p •     The last three electrons
                    3s            go into the 3p orbitals.
                         2p •     They each go into
                    2s            separate shapes
                              •   3 unpaired electrons
                    1s        •   = 1s22s22p63s23p3
Exceptional Electron
  Configurations
       Orbitals fill in order
 Lowest    energy to higher energy.
 Adding electrons can change the
  energy of the orbital.
 Half filled orbitals have a lower
  energy.
 Makes them more stable.
 Changes the filling order
     Write these electron
       configurations
 Titanium - 22 electrons
  1s22s22p63s23p64s23d2
 Vanadium - 23 electrons
  1s22s22p63s23p64s23d3
 Chromium - 24 electrons
  1s22s22p63s23p64s23d4 expected
  But this is wrong!!
      Chromium is actually:
 1s22s22p63s23p64s13d5
 Why?
 This gives us two half filled orbitals.
 Slightly lower in energy.
 The same principal applies to
  copper.
        Copper’s electron
          configuration
 Copper    has 29 electrons so we
  expect: 1s22s22p63s23p64s23d9
 But the actual configuration is:
 1s22s22p63s23p64s13d10
 This gives one filled orbital and one
  half filled orbital.
 Remember these exceptions: d4, d9
          Section 13.3
     Chemical Physics and the
    Quantum Mechanical Model
   OBJECTIVES:
    Calculate the wavelength,
     frequency, or energy of light, given
     two of these values.
          Section 13.3
     Chemical Physics and the
    Quantum Mechanical Model
   OBJECTIVES:
    Explain the origin of the atomic
     emission spectrum of an element.
               Light
 The  study of light led to the
  development of the quantum
  mechanical model.
 Light is a kind of electromagnetic
  radiation.
 Electromagnetic radiation includes
  many kinds of waves
 All move at 3.00 x 108 m/s = c
          Parts of a wave
  Crest
             Wavelength

                            Amplitude
Origin


          Trough
         Parts of Wave - p.372
   Origin - the base line of the energy.
   Crest - high point on a wave
   Trough - Low point on a wave
   Amplitude - distance from origin to trough (-)
    or crest (+)
   Wavelength - distance from crest to crest
   Wavelength is abbreviated by the Greek letter
    lambda = l
             Frequency
 The number of waves that pass a
  given point per second.
 Units: cycles/sec or hertz (Hz or sec-1)
 Abbreviated by Greek letter nu = n


               c = ln
    Frequency and wavelength
 Are inversely related
 As one goes up the other goes down.
 Different frequencies of light are
  different colors of light.
 There are a wide variety of frequencies
 The whole range is called a spectrum,
  Fig. 13.10, page 373
  Low                               High
  energy                            energy

 Radio Micro Infrared    Ultra- X-     Gamma
 waves waves .           violet Rays Rays
Low                                High
Frequency                          Frequency
Long                             Short
Wavelength                       Wavelength
                 Visible Light
                 Prism
 White light is
  made up of all the
  colors of the
  visible spectrum.
 Passing it through
  a prism separates
  it.
      If the light is not white
 By heating a gas
  with electricity we
  can get it to give
  off colors.
 Passing this light
  through a prism
  does something
  different.
         Atomic Spectrum
 Each element
  gives off its own
  characteristic
  colors.
 Can be used to
  identify the atom.
 How we know
  what stars are
  made of.
• These are called
  discontinuous
  spectra, or line
  spectra
• unique to each
  element.
• These are
  emission spectra
• The light is emitted
  given off
• Sample 13-2 p.375
        Light is a Particle
 Energy   is quantized.
 Light is energy
 Light must be quantized
 These smallest pieces of light are
  called photons.
 Photoelectric effect?
 Energy & frequency: directly related.
     Energy and frequency
E  =hxn
 E is the energy of the photon
 n is the frequency
 h is Planck’s constant
 h = 6.6262 x 10 -34 Joules x sec.
 joule is the metric unit of Energy
The Math in Chapter 13

             2  equations so
               far:
              c = ln
              E = hn
              Know these!
             Examples
 What  is the wavelength of blue light
  with a frequency of 8.3 x 1015 hz?
 What is the frequency of red light
  with a wavelength of 4.2 x 10 -5 m?

 What is the energy of a photon of
  each of the above?
Explanation of atomic spectra
 When   we write electron
  configurations, we are writing the
  lowest energy.
 The energy level, and where the
  electron starts from, is called it’s
  ground state- the lowest energy
  level.
      Changing the energy
 Let’s   look at a hydrogen atom
         Changing the energy
   Heat or electricity or light can move the
    electron up energy levels (“excited”)
         Changing the energy
   As the electron falls back to ground
    state, it gives the energy back as light
      Changing the energy
 May fall down in steps
 Each with a different energy
    Ultraviolet    Visible   Infrared
 Further they fall, more energy, higher
  frequency.
 This is simplified
 The orbitals also have different
  energies inside energy levels
 All the electrons can move around.
            What is light?
 Light is a particle - it comes in chunks.
 Light is a wave- we can measure its
  wavelength and it behaves as a wave
                           2
 If we combine E=mc , c=ln, E = 1/2
  mv2 and E = hn
 We can get: l = h/mv
 called de Broglie’s equation
 Calculates the wavelength of a particle.
         Sample problem
 What  is the approximate mass of a
 particle having a wavelength of 10-7
 meters, and a speed of 1 m/s?
 Use l = h/mv
            = 6.6 x 10-27
 (Note: 1 J = N x m; 1 N = 1 kg x
   m/s2
          Matter is a Wave
 Does  not seem to apply to large
  objects
 Things bigger than an atom
 A baseball has a wavelength of
  about 10-32 m when moving 30 m/s
 Anelectron at the same speed has
 a wavelength of 10-3 cm
 Big   enough to measure.
 The chemical physics of the
        very small
 Quantum    mechanics explains how
  the very small behaves.
 Classic physics is what you get
  when you add up the effects of
  millions of packages.
 Quantum mechanics is based on
  probability
       Heisenberg Uncertainty
             Principle
 Itis impossible to know exactly the
  location and velocity of a particle.
 The better we know one, the less
  we know the other.
 Measuring changes the properties.
 Instead, analyze interactions with
  other particles
  More obvious observations
     with the very small
 To measure where a electron is, we
  use light.
 But the light moves the electron
 And hitting the electron changes the
  frequency of the light.
  Before                After
                     Photon
    Photon           changes
                     wavelength




Moving                  Electron
Electron                    Changes
                       velocity
           Fig. 13.19, p. 382

								
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