Average Atomic Mass of “Candyum”
Chapter 3, Section 3
pp. 77-87
Name: _________________
Period: ________________
Date: _________________
Objective: To simulate how the average atomic mass of each element on the periodic table is determined, the average atomic
mass of an imaginary element, “candyum,” will be calculated from masses of candy pieces. These weighted average
calculations will then be applied to calculate the average atomic mass of real elements, and compared to the periodic table.
Background Information
1. The atomic mass of each element on the periodic table is/is not (circle one) a whole number like the mass number.
2. This is because the atomic mass of each element is a _________________ average of all occurring isotopes of that element.
3. Weighted averages are often used to calculate a student’s grade in a class (i.e. 50% tests, 25% quizzes, 25% homework). If
Polly Perfect has a 90% test average, an 85% quiz average, and a 95% homework average, determine her overall class grade.
_______________________________________________________________________________________________________
4. Let’s look at carbon. It has three naturally occurring isotopes: _______________, ________________, & _______________.
5. For a given sample of carbon (i.e. charcoal), the majority (98.93 %) of the carbon atoms of that sample are carbon-12
isotopes ( ___ protons & ___ neutrons).
6. The standard scientists use to compare units of atomic mass is the carbon-12 isotope. It has been assigned a mass of 12
amu ( ____________ ___________ ______________ ). By definition, 1 amu is exactly 1/12 the mass of a carbon-12 atom.
7. 1.07 % of carbon atoms are carbon-13 isotopes ( ___ protons & ____ neutrons), each with a mass of 13.003355 amu.
8. An even smaller portion (∼one part per trillion) are carbon-14 isotopes ( ____ protons & ____ neutrons). (14.003241 amu)
9. Perform a weighted average calculation similar to Polly’s in #3 above to determine the average atomic mass of carbon.
_______________________________________________________________________________________________________
10. Does this match the average atomic mass shown on the periodic table? __________
Directions
1.
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6.
7.
Obtain a sample of “candyum”.
Do not eat the sample, as it has been handled by many students, and your teacher has purchased some for you to eat.
Carefully separate the three isotopes (m&ms, Skittles, and Reese’s Pieces) into piles.
Count the number of each isotope and record in Table 1.
Measure the total mass of all m&ms, and record in Table 1 below. Repeat for Skittles and Reese’s Pieces.
For the remaining quantities in Table 1, use Formulas 1-5 given below.
Calculate the required quantities for the m & m isotope in the Calculations section below Table 1.
Formulas
1. Average Mass:
total mass of isotope
number of particles of isotope
2. Percent Abundance:
number of particles of isotope
x100
total number of particles
3. Relative Abundance:
percent abundance (calc. #2)
100
4. Relative Mass (rm): relative abundance of isotope (calc. #3) x average mass of isotope
5. Average Atomic Mass: rm m&ms + rm Skittles + rm Reese’s Pieces
Table 1: Isotopes Used in the Average Atomic Mass Calculation of Element “Candyum”.
Isotopes of
Element Candyum
m&ms
Skittles
Reese’s Pieces
Totals
Number
Total Mass (g)
Average Mass (g)
Percent Abundance (%)
Relative Abundance
Relative Mass (amu)
Average Atomic Mass of
Candyum: ___________
Calculations (Show all work, round correctly, and include units).
1. Calculate the average mass of an m & m isotope.
2. Calculate the percent abundance of an m & m isotope.
3. Calculate the relative abundance of an m & m isotope.
4. Calculate the relative mass of an m & m isotope.
5. Calculate the average atomic mass of candyum.
Questions
1. Which isotope of candyum do you have the most of? ___________________
2. Which quantity from Table 1 tell you this? ______________________________________
3. The average atomic mass of candyum should be closest in value to the average mass of which isotope of candyum?
(choose from: m & ms, Skittles, or Reese’s Pieces) _______________________ Explain: _________________________
_______________________________________________________________________________________________
4. Naturally occurring copper consists of 69.15% copper-63, that has atomic mass of 62.929601 amu, and 30.85% copper-65,
which has an atomic mass of 64.927794 amu. Calculate the average atomic mass, and round to two decimal places (remember
to convert percent abundance to relative abundance). Compare to the average atomic mass of copper on the periodic table.
5. Calculate the average atomic mass of argon to two decimal places given the following information: argon-36 (35.97 amu,
0.337%), argon-38 (37.96 amu, 0.063%), argon-40 (39.96 amu, 99.600%).
Compare to the atomic mass of argon on the
periodic table.
6. Create a Reference Sheet for your notebook (can be used on quizzes and tests) that includes the 5 formulas given in this
worksheet for determining a weighted average/average atomic mass. You may include an example on your Reference sheet.