Unit 3 - Atomic Structure and Electron Theory

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					Unit 3 - Atomic Structure and
       Electron Theory



                    Part I
             Chapter 4 pp. 83-91
           Chapter 12 pp. 323 - 328
The Atomists
     • “Atomists” - thought that
       matter was composed of
       tiny indivisible particles
       called atoms
     • Democritus - first
       suggested the idea of
       atoms
     • Felt that atoms were
       invisible and
       indestructible
       fundamental units of
       matter
      Fundamental Problem
• Democritus had a good idea, but did
  not prove his theory with hard facts
• Early Greek and Roman philosophers
  often theorized with little or no
  experimentation
• Not very useful for scientific
  discussion...
                John Dalton
• English school teacher
• Performed
  experiments to
  determine the ratio
  elements combine in
  chemical reactions
• Created “The Atomic
  Theory”
         Postulates of Dalton’s
           “Atomic Theory”
1) All elements are composed of substances of
   submicroscopic indivisible particles called atoms
2) Atoms of the same element are identical. Each
   atom of an element is different from another
   element
3) Atoms of different elements can physically mix or
   chemically combine in whole-number ratios to
   form compounds
4)Atoms of one element can never change into atoms
   of another element
                  Atoms
• The smallest particle of an element that
  retains the properties of that element
• Consider the Aluminum foil experiment...
  Many of you found anywhere from 1 x 104
  to 1 x 106 atoms of Al in the width of the
  foil.
• Imagine how many atoms of Aluminum
  would be in a soda can...
       Subatomic Particles
• Most of Dalton’s Atomic Theory is
  excepted today.
• However, there is one revision…
• Atoms can be broken down into even
  smaller particles
• Electrons, Protons, and Neutrons
Sir J.J. Thomson
        • 1897 - discovered
          electrons
        • Used a cathode ray
          tube to separate
          particles of different
          charges
  Cathode Ray Tube Experiment
• Found that cathode
  rays were attracted to
  metal plates that carry
  a positive charge
• Similar to magnets
• The rays were repelled
  by plates with a
  negative charge
                  Electrons




• Thomson was able to prove that all cathode rays
  are composed of e-
• Realized that electrons must be part of atoms
• Found that the electron carries one unit of
  negative charge
• 1/2000 the mass of H
                 Cathode Ray
                             Ray is bending due to negatively
• Each side of the tube      charged item here
  has a metal disc that is
  either positive or
  negative.
• The ray is negatively
  charged and moves
  from the cathode
  toward the anode
  (pos.)
                       Cathode                           Anode
             Robert Millikan
• Approximately calculated
  the charge and mass of an
  electron
• An electron has a mass of
  1/1840 the mass of a
  Hydrogen atom.
• In terms of mass that is
  9.11 x 10–28 g
                              Millikan with
                              Albert Einstein
        Protons and Neutrons
• Protons
  – Positively charged particle
  – Has a mass of 1 amu (atomic mass unit) or
    1.67 x 10-24 g
• Neutrons
  – Neutral or no charge
  – Also has a mass of 1 amu (atomic mass unit) or
    1.67 x 10-24 g
      Ernest Rutherford and the
        Gold Foil Experiment
• 1911 – Manchester,
  England
• Rutherford initially
  believed that the atom was
  hollow.
• Tested his theory of atomic
  structure by creating an
  experiment that proved
  him incorrect.
• He got a big surprise!
The Gold Foil Experiment
The Gold Foil Experiment
           • Rutherford’s
             experiment proves that
             the atom has a
             nucleus.
           • He finds that some of
             the alpha particles
             deflect!
           • Let’s take a look at
             this on a smaller
             scale…
         Rutherford’s Surprise
• Rutherford was so
  amazed by the outcome
  of the experiment he was
  quoted as saying,
  “It was almost as if you
  fired a 15-inch shell into
  a piece of tissue paper
  and it came back and hit
  you.”
              The Nucleus
• The central core of an atom, composed of
  protons and neutrons.
• Very dense
• Occupies a very small space compared to
  the rest of the atom
• Most of the small alpha particles fired by
  Rutherford were able to pass through the
  atoms
Just how big is the nucleus?
              • Most of the mass of an
                atom is located in a very
                small space.
              • In fact it takes up about
                1/10,000 of the radius of
                the whole atom.
              • If we could compare an
                atom to a football
                stadium, the nucleus
                would be the size of a
                marble in the middle of
                the stadium field.
What about the rest of the atom?
• Electrons are located in the space around
  the atom, but were so tiny they did not crash
  into the alpha particles.
• Rutherford realized that because most of the
  alpha particles passed through the gold foil,
  much of the atom must be empty space.
          Models of the Atom
• After discovering the electron, Sir J.J. Thomson
  realized that the atom was not a hard ball of
  matter, as was previously thought.
• Up to this point, atoms were considered indivisible
  using Dalton’s Atomic Theory.
• Thomson reasoned that electrons were located in
  the atom in a kind of “Plumb Pudding”
     The Plumb Pudding Model
• Thomson believed that
  the atom was made up
  of electrons that were
  “stuck” into a kind of
  positively charged
  lump of protons
• This model explained
  some of the electrical
  properties of the atom,
  but nothing really
  about it’s arrangement
Rutherford’s Model
        • After his gold foil
          experiment Rutherford
          proposed that electrons
          surrounded a dense
          nucleus that was
          actually in the center of
          the atom.
        • The rest of the atom was
          empty space
        • Later, it was shown that
          neutrons also composed
          the nucleus
           Problems with
         Rutherford’s Model
• The Rutherford model was one of the best
  in it’s time.
• However, scientists realized that positively
  and negatively charged particles would tend
  to be attracted to one another.
• What kept the particles from crashing into
  each other in a “magnetic” attraction?
          The Bohr Model
• In 1913, one of Rutherford’s students, a
  physicist named Niels Bohr, proposed that
  electrons were arranged in circular paths
  around the nucleus.
• Also known as the “Planetary Model” Bohr
  proposed that the electrons orbited around
  the nucleus like planets around the sun.
• Each electron had a fixed amount of energy
  that kept it in a path around the nucleus.
       The Bohr Model cont’d.
• Each electron moved
  within an energy level.
• Energy levels were
  regions around the
  nucleus where an
  electron was likely to
  be found.
               Energy Levels
• Bohr compared energy levels to rungs on a ladder.
• The lowest energy levels were at the bottom of the
  ladder.
• Electrons could move up and down the energy
  levels just like a person on a ladder.
• Electrons could not be located between energy
  levels, just like a person cannot stand in the empty
  space between rungs on a ladder.
• To move from one energy level to another, an
  electron had to gain or lose just the right amount
  of energy; a person climbing a ladder must move
  his/her feet just the right distance.
Ladders and Energy Levels
            • Notice that the bottom
              of the ladder, which is
              the lowest energy
              level, is closest to the
              nucleus
            • The more energy an
              electron gains, the
              higher up it moves on
              the ladder.
            • The opposite can be
              said if it moves down
              the ladder
        The “Quantum” Idea
• Scientists have labeled the amount of
  energy needed to move from one energy
  level to another as “QUANTA”
• When an electron gains or loses quanta
  (packets) of energy, the electrons move up
  or down the ladder
• Electrons that have energy are said to be
  Quantized.
       The difference between
      Energy Levels and Ladders
• One main difference between energy levels
  and ladders is that energy levels are not
  evenly spaced
• Energy levels get closer together as the
  electrons get farther away from the nucleus;
  kind of like an escalator rather than a ladder.
• It’s easy to get off at the top of an escalator.
  A person can just step off...
• Electrons also have an easier time escaping
  from an atom at the top most energy levels.
The Quantum Mechanical Model
            • 1926 - Developed by
              Erwin Schrödinger
            • Uses the quantum theory
              to write a mathematical
              equation to describe the
              location and energy of an
              electron in a hydrogen
              atom.
The Quantum Mechanical Model
• Considered the most modern description of
  the atom.
• Abstract and mathematical, this new model
  restricts the energy of electrons to specific
  values like Bohr.
• Differs from Bohr by estimating the
  location of an electron rather than assuming
  that the e- is following a given path.
               Probability
• Schrödinger uses the idea of probability to
  estimate the electron (e-) location.
• Example: If you have 3 red balls and
  1 green ball in a bag. You have a 25%
  chance of picking the green ball and a 75%
  chance of picking a red ball.
• The same principle can be applied to
  locating an electron in a given space.
     Calculating the Probability
       of finding an Electron
• The probability of
  finding an e- in a
  certain atomic volume
  is based upon the
  density of e- in the
  electron cloud.
• The higher the e-
  density, the higher the
  probability of finding
  an e-
• The opposite applies to
  low e- density
            Electron Clouds
• No identifiable line around the nucleus
• Fuzzy cloud indicates the likelihood of
  finding an electron 90% of the time.
• Example:
  – If you could put a bag around the area an e- is
    located, you would find the e- 90% of the time
    in the bag.
 The Principle Quantum Number
      & Energy Sublevels
• The P.Q.N. (n) – identifies the energy level
  that an electron is located; kind of like a
  section of a movie theater where an e- is
  located.
• As the distance from the nucleus increases,
  so also the Energy level where n = 1,2,3,4,5,
  etc,
• In each Energy Level, an e- occupies an
  energy sublevel, like a row in a section of a
  theater
                  Orbitals
• Electrons are limited by the Quantum
  Mechanical Model to cloud shapes as we
  discussed earlier.
• We call these cloud shapes orbitals
• Orbitals are the most likely location of e-
• We denote orbitals by letters instead of
  numbers (s,p,d,f)
                Orbital Types

• We will be looking at
  two different types of
  orbitals (s & p)
• s - orbital types are
  spherical shaped and
  can hold only 2 e-.
• These orbital types are
  filled first with e- on
  every energy level
                p Orbitals
• p – orbitals are filled after s – orbitals and can
  hold up to 8 e-
• There are many shapes for the p-orbital
• We call these shapes dumbbells
      Basic Nomenclature for
         Electron Theory
• To understand where e- are located we write
  e- configurations using a special
  nomenclature.
• Example
                2s 2           Number of e-

       P.Q.N.

                Sublevel and
                orbital type
We can use the Periodic Table

				
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posted:3/29/2013
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