1. Atomic Theory
   1. Postulates
      1. Dalton's Three Modified Postulates
         1. An element composed of tiny particles is called atoms. all atoms of a given element show the same chemical properties.
         2. Atoms of different elements have different properties. In an ordinary chemical reaction, no atom of any element disappears of is changed into an atom of another element.
         3. Compounds are formed when atoms of two or more elements combine. In a given compound, the relative numbers of atoms of each kind are definite and constant whole number ratios.
         4. Chemical reactions involve reorganization of the atoms-changes in the way they are bound together.

      2. Dalton's work was based on 3 laws of chemistry
         1. Law of Conservation of Mass (Antoine Lavosier, 1789),

            Based on second postulate: atoms are neither created or destroyed (under normal chemical reactions).

         2. Law of Definite Proportion, (Proust's Law),

            Based on third postulate, atom ratio is fixed, so mass must be constant.

         3. Law of Multiple Proportions (Dalton),

            Applies where two elements, A and B, form more than one compound.

            
            

2. Structure of Atoms
   1. Components of the atom

      Name	Mass(amu)	Charge	Location
      proton (p+)	1	+1	nucleus
      neutron (n0)	1	0	nucleus
      electron (e-)	0.0005	-1	outside the nucleus

      1. J.J. Thomson (1879), cathode ray tube: proved the existence of the electron by studying the deflection of cathode rays by electric and magnetic fields.

         

         m/e = 5.69x10^{-9 g}/_coulomb

         found the mass to charge ratio.

      2. Millikan (1909), oil drop experiment: determined the charge on the electron
         1.60x10^-19 coulomb

         

         Combine Thomsons and Millikans work and obtain the mass of a single electron
         m = (1.60x10^-19 coulomb)(5.69x10^{-9 g}/_coulomb) = 9.11x10^-28 g

         [1 p⁺ is 1.67x10^-24 g, therefor (1.67x10^-24g)/(9.11x10^-28 g) = 1833 about 200 times larger or 1/2000 or .005]

      3. Rutherford (1911), Gold foil experiment

         

         carried out by Johannes Geiger and Ernest Marsden.

      4. James Chadwick
         described work of Irene Joliet-Curie as neutron emission rather than a new radioactive particle.
      5. Niels Bohr-Spectroscopy working with the magnetic spectrum wanted to describe the exact location of the electron. The absorption/emission lines are the "finger prints" of the elements.
         C=v and Plank's equation E=hv

   2. Atomic number:
      atomic number = number of protons in the nucleus = the number of electrons.

   3. Mass number (Atomic Mass):

      Mass number = #(p⁺ + n⁰) = number of protons plus number of neutrons

      Atoms of the same element can differ in mass number, isotopes.

      protuim (light hydrogen), deuterium, tritium
      U235 and U238
      Symbol ²³⁵₉₂U and ²³⁸₉₂U

   4. Molecules (Compound) and Ions
      1. Compound:
      2. Ion:
         Ionic compounds will contain
         + ions (cations)
         - ions (anions)
         in a ratio that maintains electrical neutrality.

   5. Masses of atoms
      1. Gives the relative mass of an atom. Based on C-12 scale (started in 1961)
      2. Periodic table is made up of isotopic masses. These are "weighted averages" of the isotopes.

         To calculate atomic mass you must know two things: the masses of the individual isotopes and the percentages that these isotopes exist in nature.

         Ne-20	20.00	90.92%
         Ne-21	21.00	0.26%
         Ne-22	22.00	8.82%

         
         Atomic Mass = Ne = 20.00(0.9092) + 21.00(0.0026) + 22.00(0.0882) = 20.18

This page was made by Erik Epp.