ATOMS, MOLECULES, and IONS

1. Fundamental Laws of Matter

• Law of conservation of Matter - there is no detectable chnage in mass in an ordinary reaction
• Law of  Definite Proportion - a compound always contain the same element in the same proportion by mass
• Law of Multiple Proportions - if two elements can combine to form more than on type of compound, the masses of one element that combine with a fixed mass of the other element are in ratios of small whole numbers.
   

      2.The Atom
               Atom – the smallest particle  of an element that can enter into chemical  reaction.
           
        A.     Dalton’s Atomic Theory - 1808

•  An element is composed of tiny particles called atoms. All atoms of a given element show the same chemical properties. Atoms of different elements show different properties.
•  In an ordinary chemical reaction, atoms move from one substance to another, but no atom of any element disappears or is changed into an atom of another element.
•   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.

       B.      THE ELECTRON


          Cathode Ray Tube experiments were performed in the late 1800’s & early 1900’s.

•  A cathode ray tube consisted of two electrodes sealed in a glass tube containing a gas at very low pressure.
• When  voltage is applied to the cathode [negative electrode] a glow [called cathode rays] is emitted and travels to the anode [positive electrode]; hence these rays must be negatively charged.

          J.J. Thomson

•  modified the cathode ray tube experiments in 1897 by adding two adjustable voltage electrodes in order to study the amount at which the cathode ray beam was deflected by additional electric field.
• Thomson used his modification to measure the charge to mass ratio of electrons.
         ^{Charge to mass ratio:      e/m = -1.75881 x 108 coulomb/g of e-}

•  The cathode rays were named electrons.       

           Robert A. Millikan 

•    Won the 1st American Nobel Prize in 1923 for his famous oil-drop experiment. In 1909 Millikan determined that the charge on a single electron is equal to 1.60218 x 10^-19
•  Using Thomson’s charge to mass ratio, the mass of one electron can be  calculated

                                   e/m  = -1.76 x 108 coulomb
                                   e  = -1.6022  x 10-19 coulomb
                                   Thus  m = 9.10 x 10-28 g
  
       C.      THE PROTON AND THE NUCLEUS
        

         Ernest Rutherford directed Hans Geiger and Ernst Marsden’s experiment in 1910.

•   particle scattering from thin Au foils
• provided the basic picture of the atom’s structure
• In 1912 Rutherford decoded the particle scattering information.
• Explanation involved a nuclear atom with electrons surrounding the positively-charged nucleus
• The positively-charged particles in the nucleus are called ‘protons’
  

           Rutherford’s major conclusions from the particle scattering experiment

• The atom is mostly empty space.
• It contains a very small, dense center called the nucleus.
• Nearly all of the atom’s mass is in the nucleus.
• The nuclear diameter is 1/10,000 to 1/100,000 times less than atom’s radius.
  

      D.      THE NEUTRON

          James Chadwick in 1932 analyzed the results of particle scattering on thin Be films.  He recognized  existence of massive neutral particles which he called neutrons.

              Atomic Number [Z]

• The atomic number is equal to the number of protons in the nucleus
• On the periodic table, Z, is the uppermost number in each element’s box.
• In 1913, H.G.J. Moseley realized that the atomic number determines the element.
• The elements differ from each other by the number of protons in the nucleus.
• The number of electrons in a neutral atom is also equal to the atomic number.
  

              Mass Number [A]
     
          The mass number is the sum of the number of protons and neutrons.

              Mass number = number of protons + number of neutrons
              Mass number = atomic number  + number of neutrons


       E.        ISOTOPES

•  Isotopes are atoms of the same element but with different number of neutrons.
•  Isotopes have different masses but the same atomic numbers .
   

         Ex.

            Example: Isotopes of  hydrogen

             1H1  or protium is the most common hydrogen isotope
                           [1 proton, no neutrons]
             1H2  or deuterium is the second most abundant hydrogen isotope
                           [1 proton, 1 neutron]
             1H3 or tritium is a radioactive hydrogen isotope
                           [1 proton, 2 neutrons]