Docstoc

ELECTRONS

Document Sample
ELECTRONS Powered By Docstoc
					LIGHT and ELECTRONS
                 Unit 6
             Chemistry
               Langley
LIGHT and its PROPERTIES
 Pre-1900
   Issac Newton explained light and its behavior
    by assuming light moved in waves
 1900 and beyond
   Experimental evidence began to convince
    scientists that that light consists of particles
    (after the 1902 experiment of Max Planck)
 1905-Einstein
   Dual Wave Particle Theory
LIGHT and its PROPERTIES
 Wavelength: distance between two points on
  two adjacent waves, symbol is l (Greek symbol
  for lamda)
 Frequency: number of waves that pass a given
  point in a given amount of time, symbol is n
  (Greek symbol nu). Units for frequency are
  cycles per second which SI speaking is a Hertz,
  Hz (Hz is also a reciprocal seconds-1).
LIGHT and its PROPERTIES

  The frequency and wavelength of light are inversely
   proportional to each other.
    As the wavelength of light increases, the frequency
      decreases
    As the wavelength of light decreases, the frequency
      increases
  Amplitude: Wave’s height from zero to crest or wave’s
   height from zero to trough (can be positive or negative)
  A complete wave cycle starts at zero goes through its
   highest point, back through zero, reaches its lowest
   point, and back to zero again.
    One wave cycle starts at zero and has one crest
      and one trough
LIGHT and its PROPERTIES

 According to the Wave Model, light
  consists of electromagnetic waves
   Electromagnetic radiation: light moving in
    waves through space
     Radio waves, microwaves, infrared waves, visible
      light, ultraviolet waves, X-rays, and gamma
      raysElectromagnetic spectrum
   Speed of light: depending on the wavlength
    and frequency, speed of light changes
     C = ln
     Speed of light in a vacuum = 3.0 x 108 m/s
SPEED of LIGHT PROBLEMS
 EXAMPLE 1:
   Determine the speed of light if the wavelength
    is 3.5 x 10-9 m/s and the frequency is 3.5 Hz.
SPEED of LIGHT PROBLEMS
 EXAMPLE 2:
   If light has a speed of 5.6 x 103 m/s and a
    frequency of 2.3 Hz, what is the wavelength.
SPEED of LIGHT PROBLEMS
 EXAMPLE 3:
   What is the wavelength of radiation with a
    frequency of 1.5 x 1013 Hz? Does this radiation
    have a longer or shorter wavelength than red
    light?
SPEED of LIGHT PROBLEMS
 EXAMPLE 4:
   What frequency is radiation with a wavelength
    of 5.00 x 10-8 m? In what region of the
    electromagnetic spectrum is this radiation?
PHOTOELECTRIC EFFECT
(supporting work for Atomic Spectra)
 The photoelectric effect is a quantum
  electronic phenomenon in which electrons are
  emitted from matter after the absorption of
  energy from electromagnetic radiation such as
  x-rays or visible light. The emitted electrons can
  be referred to as photoelectrons in this context.
  {Wikipedia.org}
PHOTOELECTRIC EFFECT (supporting
work for Atomic Spectra)
Expected: Since all light is energy moving
 in waves, all colors of light should knock
 electrons off a metal
  Shine different color lights on a metal
  Measure the number of electrons knocked off
   the metal
  Found that no electrons were knocked off when
   light was below a certain frequency
MAX PLANCK
(his work used in Atomic Spectra)

  German Physicists, founder of quantum theory
  Studied the way light came off hot objects
   (diffusion of hydrogen through heated
   platinum)
  Concluded that light comes off in small burst of
   particles, NOT WAVES
  Quantum-minimum amount of energy that can
   be lost or gained by an atom
  To calculate quantum/energy: E = hn
    E = energy of the photon
    h = Planck’s constant
    n = frequency of the incident photon
ATOMIC SPECTRA

 As atoms absorb energy, electrons move
  into higher energy levels. When the
  atoms release energy (lose the energy),
  the electron return to the lower energy
  levels.
 The frequencies of light emitted by an
  element separate to give the atomic
  emission spectrum of the element
   No two elements have the same emission
    spectrum
ATOMIC SPECTRA

Atomic line spectra and its existence was
 known before Bohr’s atomic model of
 hydrogen was produced. What Bohr did
 was explain why hydrogen had the specific
 frequencies it had, why it “produced/broke
 down” into the colors it did; it predicted the
 values that agreed with the experiements.
 ATOMIC SPECTRA

 Hydrogen Atom Line Emission Spectrum
EXPECTED:
Continuous spectrum of
light to be given off. (Since
e- are moving around
nucleus randomly and using
different levels of energy.)
ACTUAL:
Current passed through tube with
Hydrogen gas.
Pink light is given off.
Light passed through spectrum.
Found only specific frequencies
of light given off.
ATOMIC SPECTRA
Lowest possible energy of the electron is
 referred to as its ground state
  Normal location of an electron
Electrons circle the nucleus in specific
 orbits
If an electron absorbs energy, moves up
 an energy level (absorption)
If an electron gives off energy, moves
 down an energy level (emission)
QUANTUM MECHANICS

 EINSTEIN, AGAIN!!!!!!!!!!!!!!!!
   Debate between whether light is waves or
    particles
   Einstein creates dual waves particle theory
    (1905)
   Light is small particles (photons) that move in
    wave shapes
   Thought electrons moved around the nucleus
    in wave shapes (since electrson are small
    particles like photons)
QUANTUM MECHANICS
 Louis de Broglie: Given that light behaves as waves and
  particles, can particles of matter behave as waves?
    Referred to the wavelike behavior of particles as matter
     waves
    Came up with an equation that predicts all moving objects
     have wavelike behavior:
       mv/l = h
       Thanks to experiments conducted by 2 scientists, his theory
        was proven correctNobel Prize
       Waves Waves have specific frequencies and electrons have
        specific orbits/energy levels
       Waves and electrons can both be bent (diffraction)
       Waves and electrons can both overlap and interfere with
        each other (interference)
    Creator of Wave Mechanics
QUANTUM MECHANICS

 DeBroglie’s equation combines Einstein
  and Planck’s equations
   mv/l = h
   (Anything with mass and velocity has a
    wavelength, so electrons have wavelengths)
   DeBroglie Problems:
     What is the wavelength of an electron that has a
      mass of 1.5 X 10-30 kg and a velocity of 2.5 X 104
      m/s?
QUANTUM MECHANICS

  DeBroglie Problems:
    What is the velocity of an electron with a mass of
     8.3 X 10-29 kg and a wavelength of 400 nm? (Hint:
     convert nm to m)
    What is the mass of an electron with a velocity of
     4.6 X 103 m/s and a wavelength of 5.6 X 10-2
     meters?
    What is the wavelength of an electron that has a
     mass of 2.8 X 10-31 kg and a velocity of 3.0 X 108
     m/s?
QUANTUM MECHANICS
 Heisenberg
  2 Goals in Life:
      find the location of an electron
      find the velocity of an electron
  Problem: Electrons cannot be seen under a microscope
  Only way to find an electron is to shoot a photon
   (particle of light) at the electron
  Problem: when the photon hits the electron, it knocks
   the electron off course
  So with this photon method, you can only know the
   position of an electron for a split second, but you still
   don’t know the velocity
QUANTUM MECHANICS

Heisenberg
  DeBroglie: Tries to help Heisenberg and offers
   his equation l = (mv)/h
  If you know mass and wavelength of an
   electron, equation could help you find velocity
  Problem: Equation does not show location!
  Equation method will only tell you velocity NOT
   location
QUANTUM MECHANICS

Heisenberg
  Heisenberg Uncertainty Principle: It is
   impossible to know both the position and
   velocity of an electron at the same time.
QUANTUM MECHANICS
 Schrodinger
  Working with Hydrogen atom that only has 1 electron
  Wants to find general location/area of the one electron
   in Hydrogen
  Creates quantum theory
  Quantum theory – uses math to describe the wave
   properties of an electron (frequency, wavelength, etc)
  Once he plugged his data into the quantum theory, he
   found that electrons do not travel in nice, neat orbits
   (Bohr model)
  Instead, found that electrons travel in 3D regions around
   the nucleus
QUANTUM MECHANICS

Schrodinger
  Schrodinger’s equation is used to find the
   greatest probable location/area of the Hydrogen
   atom electron (in the ground state)
QUANTUM MECHANICS
 Quantum Theory
  Ground State-normal location of an electron
  Excited State-one ring up from the normal location
  When excited electron falls back to the ground state, a
   photon is given off
  Energy of the photn is equal to the difference in energy
   between the excited state and ground state
  Hydrogen gives off specific colors because its electrons
   move from ring 2 to ring 1; Neon gives off a different
   color because its electrons move from ring 3 to ring 2
LIGHT AND ELECTRONS REVIEW
 Light was first thought to be wavelike
 Equation for the speed of light is c = ln
 Photoelectric effect challenges this because only certain
  frequencies of light could knock off electrons
 Max Planck’s experiment proved that light could be a
  particle
 Einstein’s dual wave particle theory says that light is
  ACTUALLY small particles (photons) that move in wave
  like patterns
 Equation for energy of a photon is E = hn
 Bohr found that electrons orbit the nucleus in specific
  orbitals/energy levels
LIGHT AND ELECTRONS REVIEW
 Electrons as Waves:
  1924 – Louis de Broglie asked “Could electrons have a
   dual wave particle nature like light?”
      Similarities between waves and electrons
      Waves have specific frequencies and electrons have
       specific orbits/energy levels
      Waves and electrons can both be bent (diffraction)
      Waves and electrons can both overlap and interfere with
       each other (interference)
      DeBroglie’s equation combines Einstein and Planck’s
       equations
      mv/l = h
      (Anything with mass and velocity has a wavelength, so
       electrons have wavelengths)
QUANTUM NUMBERS and ATOMIC ORBITALS

  REVIEW
   Energy levelsSpecific energies electrons can have
   Quantum of energyamount of energy required to
    move an electron from one energy level to another
    energy level
   The amount of energy an electron gains or loses in an
    atom is not always the same
   Energy levels in an atom are not equally spaced
   Higher energy levels are closer together
   Modern description of the electrons in atoms, quantum
    mechanical model, comes from the mathematical
    solutions to the Schrodinger equation
   The quantum mechanical model determines the allowed
    energies an electron can have and how likely it is to find
    the electron in various locations
QUANTUM NUMBERS and ATOMIC ORBITALS

 QUANTUM NUMBERS
   Quantum numbers are used to describe the
    location and behavior of an electron (zip code
    for electrons)
     First quantum number = Principal = n
     Second quantum number = Angular Momentum
     Third quantum number = Magnetic Quantum
     Fourth quantum number = Spin Quantum
QUANTUM NUMBERS and ATOMIC ORBITALS

  Principal (first) quantum number = n
         Main quantum number
         Describes the size of the electron cloud (the smaller the number,
          the smaller the cloud)
         ALSO, shows the distance from the nucleus, the smaller the
          number, the closer the cloud is to the nucleus
         Called energy levels or shells
         Positive integers (1,2,3,4,…)
         Symbol is n
         Each energy level has a maximum number of electrons it can hold
                n                   # Electrons
                1                       2
                2                       8
                3                      18
                4                      32
         Example: Energy level 1
                    2 electrons
                    close to the nucleus
                    small electron cloud
QUANTUM NUMBERS and ATOMIC ORBITALS

 Second Quantum Number:
   Each energy level has sublevels
   The number of sublevels is equal to n
   Example: Energy level 1 has 1 sublevel
   Sublevels are called: s,p,d,f
QUANTUM NUMBERS and ATOMIC ORBITALS

  Third Quantum Number
    Divides sublevels into orbitals
    Also tells the shape the electron is moving in
    The number of orbitals for each level is:
        S has 1
        P has 3
        D has 5
        F has 7
    The number of orbitals for an energy level is equal to n2
    Example: 2nd Energy level
        n2 = 4
        1s, 3p
    Each orbital can only hold a maximum of 2 electrons
    Shapes of orbitals:
        S is spherical
        P is peanut shaped
        D is daisy shaped
        F is unknown
QUANTUM NUMBERS and ATOMIC ORBITALS

 Fourth Quantum Number:
   Describes the electron spin
   Both electrons in an orbital are negative, so
    they repel each other and spin in opposite
    directions
   Use arrows to represent electrons
QUANTUM NUMBERS and ATOMIC ORBITALS

 Pauli Exclusion Principle:
   No two electrons in an atom can have the
    same set of 4 quantum numbers because
    electrons repel each other
   Example: 2 electrons may both be:
     in the first energy level (same first number)
     sitting in an s sublevel (same second #
     moving in a sphere shape (same third #)
     BUT one electron spins clockwise and one
     spins counter clockwise ( which means they
     have different fourth #s)
     ELECTRON CONFIGURATIONS

Example 1: Map out the quantum
 numbers for all the electrons in Hydrogen
  Find the # of electrons in hydrogen (atomic
   number will give you this number)
  What order do you fill in s, p, d, f in the rings?
      ELECTRON CONFIGURATIONS

 Diagonal RulePattern that shows the order the
  electrons fill in the orbitals: Some People Do Forget
                               Notice that the electrons do
1s                             not fill in all of the level 3
2s    2p                       first (3s, 3p, 3d) and then
3s    3p     3d                move to level 4
4s    4p     4d     4f          Instead, electrons fill in the
5s    5p     5d     5f          orbitals in the order that is
6s    6p     6d                 easiest to them (easier for
7s    7p                        an electron to fill in a 4s
                                before it fills in a 3d)
  Aufbau Principle: Electrons have to fill in the lowest
  (easiest) energy level or orbital first
ELECTRON CONFIGURATIONS
 Hund’s Rule:               Example 2: Helium
  Every orbital must get
   one electron first,
   before you double up.
   “Cookie Rule”
ELECTRON CONFIGURATIONS
 Example 3: Lithium    Example 4: Fluroine
ELECTRON CONFIGURATIONS
 Orbital Notation        Electron
  drawing out             Configuration
   configurations with     Notation/Superscript
   arrows                  Notation:
                           writing configurations
                            with superscripts to
                            represent electrons
ELECTRON CONFIGURATIONS

 Do Orbital Notation and Electron
  Configuration for the following:
     Zn
     I
     Cl
     Mg
     As
NOBLE GAS CONFIGURATIONS
 Noble Gas Configurations:
 Write out the superscript notations for:
       Neon:
       Sulfur:
       Sulfur has the same configuration as Neon plus a 3s23p4
       So, you could use the noble gas as a shortcut and write
       Sulfur’s configuration as
       [Ne] 3s23p4 OR            [Ne]
 Noble gas configuration: write the noble gas (group 18) that comes
  directly before the element in question and then add the rest of the
  configuration
 Practice:
     Write the noble gas superscript notation for the following elements.
         C                  Np
         W                  Sn
DOT DIAGRAMS

Lewis Dot Diagrams:
  Way to show the number and position of
   electrons on the outermost energy level
  Since the energy levels all overlap and cover
   one another, only the outermost energy level is
   able to bond with other elements
    The electrons involved in bonding are called the
     valence electrons (to get these electrons look at the
     column number)
DOT DIAGRAMS

Lewis Dot Diagrams:
  Chemical symbol + Number of valence
   electrons
  The rules for orbitals still apply, so no side can
   have more than two dots, and each “p” orbital
   side gets one dot, before you double up
                               p1
                                        s orbital
   p orbitals             p2
                               Xs
                               p3
DOT DIAGRAMS

Noble gases have a full valence
There are no empty spaces so the
 element does not need any more electrons
Stable octet – 8 electrons in the valence
 so the element does not want to bond (this
 means it is stable)
Only the noble gases have a stable octet
DOT DIAGRAMS

Practice: Write the noble gas superscript
 notation and then draw the dot diagram for
 the following:
  V
  Br
  Al
  K

				
DOCUMENT INFO
Shared By:
Categories:
Tags:
Stats:
views:9
posted:2/8/2012
language:English
pages:46