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Atomic history and quantum numbers

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					SCH 4U0
Unit 1: Structure & Properties
Note: Development of the Atom and Quantum Numbers
                                    Development of the Atom

                                                                      FIRE        EARTH
Ancient Greeks
450 BC Empedocles
        - 4 elements and properties
400BC Democritis                                                     AIR           WATER
        - keep dividing in half , "atomos"
Aristotle – agreed with Empedocles, had 2000 yrs. Of influence on scientific community

Bottom Line: Great Thinkers, no experiments to prove right or wrong

Alchemists
    found uses of elements, not explaining their behaviour
    thought they could change iron into gold
    didn’t follow scientific method

Bottom Line: lots of experiments and equipment

Early Scientists (carefully controlled experiments)
Robert Boyle – not just 4 elements
Joseph Priestly – discovered oxygen (discovered rust)
Antoine Lavoiser – accurate measurements (made first balance)
Joseph Proust – Law conservation of mass and constant composition

Bottom Line: Scientific Method created, start to explain behaviour

JOHN DALTON 1805 (English teacher)
   combined several ideas and experiments into one theory that can be used to predict the outcome of
     other experiments
   his experiments supported:
     - The Law Conservation of Mass
     - The Law of Definite Composition (elements combine in specific mass ratios)
     - The Law of multiple Proportions (more than one combining capacity that results in more than one
         mass ratio)

Model: Indivisible billiard ball



Bottom Line: Combined great ideas into one theory to explain experiments

J.J. THOMSON
    - new studies indicated that electrical charges were apart of matter (Faraday & Crookes)
    - atoms are composed of electrons embedded in a positively charged sphere

Model: raisin bun



Bottom Line: Electricity
SCH 4U0
Unit 1: Structure & Properties
Note: Development of the Atom and Quantum Numbers
ERNEST RUTHERFORD
   -   studied radiation which allowed for the gold foil experiment
   -   alpha particles (He nucleus) were shot at a thin piece of gold foil

prediction:                                  observation:




   -   severe deflection indicated that a dense core must be present
   -   core is positive therefore there must be a strong nuclear attraction keeping the positive charges
       together
   -   most went straight through therefore, most of atom is empty space (electron cloud)
   -   found H+ is smallest positive charge = proton
   -   discovered proton had equal & opposite charge as an electron and 1836 x more mass
   -   worked with Chadwick (mass spectrometry pg. 165) who discovered another particle was
       contributing to the mass of an atom that is not charged = neutron
   -   isotopes are explained

Model: positive with electron cloud




The Problem With Rutherford:
   - Black Hole Catastrophe: Why don't electrons spin into the nucleus?
   - if they are always changing direction then they are always accelerating
   - if always accelerating then always giving off energy as light
   - if give off energy then they should eventually fall into the nucleus and cause a black whole effect
      (intense gravity pull on matter, even light)

Contributions of Bohr, Einstein, Plank:

Planck:
   - -energy is not continuous, it is in levels or "quanta" (Max Plank, coin analogy)

Einstein:
   - explained photoelectric effect (metal give off electrons when bombarded with heat)
   - light consists of photons which are essentially a quantum of energy, different wavelengths represent
       different energies E = hf and c = f 
   - so E = h (c/)
   - this energy can be transferred to electrons so they can break free of the atom
SCH 4U0
Unit 1: Structure & Properties
Note: Development of the Atom and Quantum Numbers
Bohr:
  - EMR, line spectrum of elements (energy is given off as light)
  - energy (light) is given off as the electron falls to a lower energy level
  - ground state = lowest energy level, excited state = when it has absorbed energy & is in a higher
      level
  - different colours = different wavelengths = different energies, so electrons must be in different
      energy levels
  - electrons must absorb a certain quantum of energy to move up level (like stairs)
  - periods are explained by the number of electrons that fit in an energy level
  - line spectra is different for each element (like a fingerprint)

Model: - protons and neutrons in nucleus with energy levels




Bottom Line: Electrons orbit in energy levels that have a specific quantity of energy.


PRESENT DAY MODEL
  - Bohr model only supported by hydrogen (single electron systems)
  - Why is there 2e- in the first and 8e- in the second?

Louis de Broglie 1923
   - made the hypothesis that an electron can act like a wave and particle and that each has a specific
       quantum of energy
   - proved this mathematically, proven experimentally later by G.P. Thomson

ERWIN SCHRODINGER:
     - Heisenberg Uncertainty Principle – can't know exact momentum and position of the electron at
        the same time
     - to measure one you will change the other
     - to measure the position you have to hit it with a smaller wave  increased its energy so now it
        has a different position (cat analogy)
     - this led to probability math (electron is in this general area 95% of the time)
     - electron is viewed more as a wave than a particle therefore can be determined using
        mathematical equations
     - combined probability mathematics to come up with the "Quantum Model"

   Quantum Model:
     - electrons are assigned specific number (like a postal code) derived by mathematical equations
         (more than one possible number i.e. quadratic equation)
     - the number is the general orbital (not an orbit b/c not a specific place) where the electron is
         found (electron probability density clouds formed)
     - electrons have specific energy and orientation in space, and a spin
SCH 4U0
Unit 1: Structure & Properties
Note: Development of the Atom and Quantum Numbers
                                          Quantum Numbers

   Results in 4 Quantum Numbers For Each Electron:

   Quantum            Energy             Orbital               Value            Connection to
   Number             Connections        Connections                            Periodic Table
   Principal QN (n)   main energy        shell                 n = 1          Row
                      level
   Secondary QN       sub energy level   orbital shape         l = 0  n-1      block
   (l)
   Magnetic QN        energy in          orbital orientation   ml = -l  +l     column
   (ml)               magnetic field     in space              except for l=o
   Spin QN (ms)       individual e-      electron spin         ms = -1/2 or     Individual
                      energy                                   +1/2
                      difference

   Example He = 1s2 ___

    Bottom Line: Know how to connect quantum numbers to the shape of the periodic table.




Pauli Exclusion Principle (spins)
       - No two atoms may have all four quantum numbers alike
       - The first three quantum numbers determine a specific orbital
       - As a result, only two electrons may exist in the same orbital and these electrons must have
          opposing spins

Hund's Rule (no pairs unless must)
      - When orbitals of identical energy level are available, electrons occupy these orbitals singly and
         in parallel spins
      - Electrons carry similar charges and therefore spread out as far as possible from each other
      - Therefore, one electron in each orbital unless must pair up

Aufbau Filling Principle
      - Aufbau (German) means “to build up”
      - An imaginary process of placing electrons in their ground state configurations
      - Therefore, fill lower levels first
SCH 4U0
Unit 1: Structure & Properties
Note: Development of the Atom and Quantum Numbers

Why is it that the 4s orbital is filled before the 3d orbital?




    -   As atoms become larger the main energy levels become closer together and therefore some of the
        sublevels start to overlap in energy (P. 187 of text).




    Fig: Shows relative energies of electrons in the
          different orbitals

				
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posted:11/26/2011
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