Atomic Structure Atomic Structure

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					Atomic Structure

Small-Scale Architecture
Chapter 4
Chapter 4 Objectives
   Understand different atomic models.
   Describe the arrangement and the
    electron capacities of the various
    principle energy levels, sublevels, and
   Place electrons in orbital notation
    according to the Aufbau principle and
    Hund’s rule
   Write out the electron configuration of
    common elements.
Chapter 4 Objectives (cont.)
   Describe the information given by each
    of the four quantum numbers.
   State the number of valence electrons
    in a given atom.
   Use electron dot structures to represent
    valence electrons.
4A- The Development of Atomic
Models: A Historical Perspective
 Models- working representations of experimental
    • Models do not give us a complete picture. Instead, they give
      a simplified picture that is more easily understood.

                          Plum Pudding
4A- (cont.)

   The Origin of the Atomic Concept: Greek Ideas
    • The Greek philosopher, Democritus used the term atomos
      from which we get the term atom.
        – Matter is discontinuous and is made of separate, definite
          particles that cannot be infinitely divided.
        – If divided enough times, there would be a particle that could not
          be further subdivided without loosing its inherent properties
    • Atom- the smallest part of an element that retains the
      properties of that element.
    • The Law of Definite Composition states that every
      compound is formed of definite particles combined in definite
        – Every molecule of water contains the same two elements,
          Hydrogen and Oxygen, in the same proportions. Their
          proportions by mass are 8.00 g of Oxygen for every 1.00 g of
4A- (cont.)
       The First Experimental Model: Dalton’s Atomic
    •     First to frame a model based on experimentation.
    •     Summary of Dalton’s Model
    1.    Elements are made of minute particles called atoms,
          which are tiny, indestructible spheres.
    2.    Atoms of different elements have unique sizes and
          properties. (Different masses)
    3.    An atom of one element cannot be changed into an atom
          of another element.
    4.    Atoms form compounds by combining with each other.
    5.    A certain compound always contains the same relative
          number and kinds of atoms.
    •     Dalton began assigning relative mass numbers to each of
          the elements. The numbers were not completely accurate.
4A- (cont.)
   Discovery of the Electron: Thomson’s Model
    • Development of the battery sparked interest in electricity.
    • Gases sealed in a glass tube (gas discharge tube) under low
      pressure conduct a current between two electrical contacts
      called cathodes.
       – Initially, removing the gas with a vacuum pump decreased the
         current. (There were fewer molecules to carry the current)
       – Removing more gas caused a gradual increase in the current
         and the tube began to glow green.
       – Further removal of the gas caused the glow to stop, but the
         current continued.
       – The current flowed across something other than gas.
       – Because the current came from the cathode it was called
         cathode rays.
4A- (cont.)
  • An English physicist, J.J. Thomson, made the following
     – Cathode rays travel in straight lines and are unaffected
       by gravity.
     – A magnet could deflect cathode rays.
           Conclusion: the cathode rays are composed of particles.
      – When a cathode ray passed between two charged plates
        the ray bent towards the positively charged plate.
           The particles must be negatively charged.
  • Cathode rays must be made of small, negatively charged
    particles moving at high speed.
  • The particles had a relatively large charge and a very small
    mass. He called the particles, electrons.
      – Atoms of every element emit these particles.
      – Mass- 9.11 X 10-31
4A- (cont.)

                Plum Pudding Model
  •   Thomson’s new atomic model:
      1. Negatively charged particles are embedded in
         a positively charged substance.
      2. The positive charge balances out the
         negative charge so that the atom is neutral.
      3. Under certain conditions electrons could be
         removed from an atom.
4A- (cont.)
   Discovery of the Proton: Rutherford’s Model
    • Ernest Rutherford’s work with radioactive alpha particles
      (large positively charged ions emitted from radioactive
      elements at high speeds) led to the discovery of protons.
        – Rutherford passed a beam of alpha particles through a thin
          gold foil. Most of the beam passed through, but a few of the
          particles were deflected.
    • Atoms must have a dense, positively charged center, the
    • He reasoned that since only a few of the particles were
      deflected the nucleus must be small.
    • The nucleus is about 1/100,000 the size of the atom.
    • Protons-the small, positively charged particles in the
        – Charge of +1, exactly opposite of an electron
        – Mass of 1.67 X 10-27 Kg, 1836 times the
          mass of an electron
4A- (cont.)

   Discovery of the Neutron: Chadwick’s Work
    • Scientists realized that atoms are much more massive than
      they expected based on the masses of the protons and
    • In 1932, an English Physicist, James Chadwick, identified
      the neutral particles present in the nucleus that account for
      the additional mass of an atom.
    • Neutrons- small neutral particles present in the nucleus of
      an atom.
        – Mass- 1.68 X 10-27 Kg
4A- (cont.)
   Energy Levels for Electrons: Bohr’s Model
    • If protons are positively charged and electrons are negatively
      charged, what keeps the electrons from being pulled into the
      nucleus and collapsing the atom?
    • It was believed that electrons moved around the nucleus and
      the momentum kept them from entering the nucleus.
    • Spectroscopy- the study of how atoms absorb and emit light
         – Continuous spectrum- light that separates into all the colors
           of a rainbow as it passes through a prism.
                The sun and incandescent bulbs
         – Line spectrum- a set of colored lines and bands
                Each element produces its own unique line spectrum.

      Continuous Spectrum
4A- (cont.)

  • Bohr developed an atomic model where electrons could exist
    only in definite energy levels outside the nucleus.
      – Low energy levels are close to the nucleus and higher energy
        levels are farther from the nucleus.
      – Electrons normally exist in the lowest possible energy level
        called the ground state.
      – When the right amount of energy is applied the electrons
        become excited and jump to a higher energy level.
      – The electrons return to a lower energy level, giving off the
        excess energy as light. This causes the line spectrum. The
        observed color tells of the particular energy level.
  • Bohr was able to predict the exact wavelength of the lines in
    the hydrogen spectrum.
4A- (cont.)
  • Principle Energy Levels- Bohr’s energy levels, a region
    around the nucleus containing a specified group of electrons
    in sublevels and orbitals.
  • Six or seven principle energy levels can be measured.
    Theoretically, more may exist.
  • Each principle energy level has a maximum number of
    electrons that it can hold.

                       Principle      Max. No. of
                     Energy Level     Electrons
                           n        2n2 (observed)

                          1               2

                          2               8

                          3              18

                          4              32

                          5            50 (32)

                          6            72 (18)

                          7             98 (8)
4A- (cont.)
   Evidence for the Quantum Model: New Physics
    • Bohr’s model works well for Hydrogen but needed
      adjustments for elements with more than one electron.
    • Heisenberg Uncertainty Principle- it is not possible to
      know both the energy (momentum) and the exact position of
      an electron at the same time.
        – All communication between an electron and an observer has to
          be mediated by photons (light particles) or other electrons.
        – Looking at an electron changes its position by way of a
          subatomic collision.
    • Bohr’s orbits were replaced with orbitals- 3-dimensional
      regions where electrons probably exist. Four dimensional if
      you include a time.
4A- (cont.)
  • Principle energy levels are still arranged around the nucleus.
      – Orbitals in which the electron’s average distance is close to the
        nucleus have low energy.
      – Orbitals in which the electron’s average distance is far from the
        nucleus have high energy.
  • Are electrons particles or waves?
      – In 1924 Louis de Broglie stated that the matter of an electron
        was not concentrated at one point but was spread out over the
        entire orbital.
      – Electrons are said to have a dual nature. They act as a particle
        and a wave.
  • Difficult concepts bring us to the realization that an infinitely
    wise creator made and designed the universe according to
    His plan.
4B- The Quantum Model: Where are the Electrons?

       • Most recent atomic model

   Sublevels and Orbitals
    – Electrons exist in principle energy levels
    – Principle energy levels are subdivided into
      sublevels-made up of one or more orbitals
       • s, p, d, f
    – S sublevel
       • Simplest of all sublevels
       • Spherical shape
       • Only one orbital holding a total of 2 electrons
4B- The Quantum Model: Sublevels and Orbitals

   – p sublevel
      • Shape resembles 3 sets of barbells
      • Three p orbitals (p1, p2, p3)

      • Holds 6 electrons
   – d sublevel
      • Five orbitals
      • 10 electrons
   – f sublevel
      • Seven orbitals
      • 14 electrons
   – Each principle energy level as specific sublevels
– First principle energy level has only an s
– Second principle energy level
  has an s & p orbitals (8 electrons)
– Third level contains s,p,&d
  orbitals (18 electrons)
– The fourth & fifth levels have
  s, p, d, & f orbitals with (32 electrons)
   • Higher principle energy levels do not contain all
     the orbitals they could because the known
     elements only have a certain number of
4B- The Quantum Model: Energies of Sublevels

   – Sublevels are ranked according to their
   – 1s 2s 2p 3s 3p 4s 3d 4p 5s 5p 6s 4f 5d 6p 7s 5f 6d 7p
   – Several of the principle energy levels
   – Diagonal Rule
          Diagonal Rule
1s                  1.   Write the energy levels top to
                    2.   Write the orbitals in s, p, d, f order.
2s   2p                  Write the same number of orbitals
                         as the energy level.
                    3.   Draw diagonal lines from the top
3s   3p   3d             right to the bottom left.

                    4.   To get the correct order,
4s   4p   4d   4f        follow the arrows!
                                   There are no known

5s   5p   5d   5f        5g?       elements that go this high!

6s   6p   6d   6f        6g?           6h?

7s   7p   7d   7f        7g?           7h?           7i?
3B- The Quantum Model: Aufbau Principle-How to
Build an Atom

   – The arrangement of electrons in an atom
     may be determined by the addition of
     electrons to a smaller atom.
   – German for “building up”
   – As the Periodic Table progresses each
     succeeding element has one additional
     proton and one additional electron.
   – Lower energy orbitals must be filled 1st.
Aufbau cont.
  – Electron Configuration-the arrangement of
     • Hydrogen’s one electron resides in the 1s sublevel.
     • Its electron configuration is 1s1
  – Orbital Notation: representation using dashes and
    arrows to show principle energy levels, and
    orbitals in an atom
     • _ represents an orbital
      represents the first electron in the orbital
      represents the second electron in the orbital
                            H 1s1
     • Examples!
Hund’s Rule

   –Hund’s rule states that
       as electrons fill a
     sublevel, all orbitals
     receive one electron
    before any receive two.
3B- The Quantum Model: Addresses for Electrons

 – Quantum numbers describe locations and
   energies of electrons
    • Each electron has 4 numbers
    • n- princile energy level (1,2,3,4,5,6,7) 2n2
    • l- type of sublevel (s=0,p=1,d=2,f=3)
    • m- the electron’s orbital 2 l +1

                                    3 Quantum Number
2 quantum number
        l                                Possibilities
                       2l   +1
      0 (s)                 1                 0
      1 (p)                 3               -1 0 1
      2 (d)                 5             -2 -1 0 1 2
3B- The Quantum Model: Addresses for Electrons

      • ms differentiates between 2 electrons in an orbital.
          – 1st one is assigned + ½
          – 2nd one is assigned – ½
      • Sample:
          – Give the values of n,l,m for the following:
              » 2p
              » n=2, l=1, m= -1,0,1
              » 3s
              » n=3, l=0, m=0
              » 4d
              » n=4, l=2, m=-2,-1,0,1,2
3C- Numbers of Atomic Particles: Things We Can Count on

       • The number of protons in an atom determines
         its identity; the number of neutrons affects its
         mass; and the number of electrons affects its
         electrical charge and interaction with other
   Atomic Mass
    – Atomic Mass Units (amu)- 1/12 the mass of one
      atom of carbon 12; 1.6606 x 10-27 kg
    – Atomic mass number- the sum of the protons
      and neutrons
       • Approximates the mass of the protons and neutrons.
    – Atomic number- number of protons in an atom,
    determines its identity
3C- Numbers of Atomic Particles: Isotopes

   Isotopes: Count those Neutrons!
    – Isotope- an atom with the same atomic number
      but different numbers of neutrons

    – Mass number- atomic number = number of neutrons
    – Isotopic Notation- specifies the exact composition
    of isotopes
         • Atom's symbol
         • Atomic number
         • Atomic Mass Number
3C- Numbers of Atomic Particles: Isotopes
   – Atomic Mass- weighted average of isotopes,
shows the average mass of an atom
            • Atomic number
            • Atomic Mass Number
   – Samples
             •   Chlorine has two isotopes Cl-35 &
                  Cl-37. Cl-35 is 75.77% abundant and Cl-
                 37 is 24.23% abundant.
             •   In a sample of 100 atoms there would be
                 75.77 atoms of Cl-35 & 24.23 atoms of Cl-
             •   75.77 X 34.969 amu = 2650 amu
             •   36.966 X 24.23 amu = 895.7 amu
3C- Numbers of Atomic Particles: Valence Electrons

   Valence Electrons: Last, but not least
    – The electrons in the outermost shell are involved
      in bonding.
        • Argon 1s2 2s2 2p6 3s2 3p6
           – 3rd Shell contains 8
           – Argon contains 8 valence electrons
        • Nickel 1s2 2s2 2p6 3s2 3p6 4s2 3d8
           – The outermost shell is the 4th shell
           – It has 2 electrons
           – Nickel contains 2 valence electrons
3C- Numbers of Atomic Particles: Electron Dot

• Electron Dot Structures