Isotope Worksheet - PowerPoint

							                  isotope
• Are atoms of the same element that have
  different masses. Due to a different number
  of neutrons in the nucleus.
                  Nuclide
• Is a general term for any isotope of any
  element.
            Atomic number
• Represented in mathematical equations by
  the letter Z in the book examples.

• Remember. The number of protons is
  represented by the atomic number, and will
  never change. It is the “fingerprint” of the
  element.
             Mass number
• Is the total number of protons and neutrons
  in the nucleus of an isotope.

• The neutron and the proton mass values are
  each approx 1.
• The periodic table reports the average
  atomic mass of the many nuclides of a
  particular atom
           Atomic mass unit
• The mass of atoms is extremely small.
• Therefore, we use relative atomic masses
  because it is easier to do calculations.
• Arbitrarily chose a standard and assigned a
  relative mass value.
• Selected the carbon –12 nuclide.
     Atomic mass unit or amu
• Carbon -12 has been assigned the relative
  value of 12 amu.
• Or 1 amu is exactly 1/12th the mass of a
  carbon-12 atom ( nuclide)
• Mass of any nuclide is compared to the
  carbon –12 nuclide
• An atom that is 4x the mass of the carbon-
  12 atom will be 4 x 12 amu = 48 amu.
                continued
• An atom that is 2/3 the mass of the carbon-
  12 atom will have a mass of 2/3 x 12 amu =
  8 amu.
   Finding average atomic mass
• Multiply the abundance factor, (%), by the
  mass number of the isotope
• Add up all the isotopes
• Divide by 100
                 example
• Neon has two isotopes Ne-20 has a mass of
  19.992 amu and occurs in nature 90% of the
  time. Ne-22 has a mass number of 21.991
  amu and occurs 10% of the time.

• ((90 x 19.992) + (10 X 21.991) ) / 100
  = 20.192 amu
      How many in a dozen?
• How many in 2 dozen
• If you have 36 shoes, how many dozen?
                     Mole
• The chemist dozen
• Instead of 12 units of something it is
      6.02 x10 23 units of something.
• 6.02 x 1023 is called Avogadro’s number
  – Or the number of particles in exactly one mole
    of a substance.
              Molar mass
• The mass of one mole of a substance.
• Units are g/mol
• Numerically equal to the atomic mass listed
  on the periodic table.
• The molar mass of an element contains one
  mole, or 6.02 x1023 atoms of that element.
               conversions
• He has a molar mass of 4.00 g/mol
• How many grams of He are there in 5
  moles?
• 4.00 g/mol x 5 mol = 20 grams

• Note: the mole unit crosses out.
                problem
• How many grams will there be in 4 moles
  of oxygen?
                  again
• How many grams are there in 4.5 moles of
  copper?




• Page 83 sample problem 3-2   1,2,3,4
         Now the other way
• Given the grams calculate the moles.
• You have 72 grams of copper, how many
  moles do you have?

• 72 grams / 63.55 g/mol = 1.13 moles of Cu

• Page 83 sample problem 3-3 1,2
How many moles is a given # of
          atoms?
• If you have 2.0 x 10 23 atoms of Ca, how
  many moles do you have.

• 2.0 x 1023 atom Ca x 1 mol Ca/6.02x 1023
  atom of Ca = 0.33 mol Ca

• Page 84 sample problem 3-4 1,2,3
   Given # atoms find the mass
• Find the mass of 5 x 10 16 atoms of copper.

• 5 x 10 16 atoms Cu x 1mole/6.02x10 23atoms
  multiplied by 63.55 g/mol = 1.31 x 10-9


• Page 85 practice problem 3-5 1,2,3