## CHEM& 140 Lab: Isotopes and Atomic Mass

## Part 1: Make Isotopes

1. What type of particles determine the identity of the atom, meaning which element it is?

2. What type of particles determine the mass number?

3. Why is mass number always a whole number?

4. One isotope of carbon (C) has exactly the same mass number as the value for the atomic mass since it was used as the definition of the atomic mass unit (amu). Which isotope of carbon is it and what is its mass?

5. What is the approximate mass of one proton? __________amu

6. What is the approximate mass of one neutron? __________amu

7. Look at 3 or 4 other atoms using the simulation. Do any of them have a whole number for their mass?

The reason why is that only one particular isotope of carbon, which you identified a couple of questions previously, is used to set (define) the atomic mass unit.

Fill in the blank with a fraction: 1 atomic mass unit (amu) = _________ mass of a carbon-12 atom

This turns out to be: 1 atomic mass unit (amu) = 1.66 x 10 – 27 kg

Click on the element oxygen in the Periodic Table in the simulation. Start with an atom that has 8 protons, 8 neutrons, and 8 electrons.

8. How does the mass (in amu) of the entire atom relate to the mass number?

9. Write the symbol for this atom of 8 protons, 8 neutrons, and 8 electrons using the notation style.

10. How many different atoms can you build for the element oxygen using the simulation? Write the symbols for all of these atoms using the notation style.

11. The different atoms that you just listed for oxygen are the ** isotopes** of oxygen. Which of the isotopes for oxygen are stable? Which are unstable? Which isotopes are abundant in nature, the stable or unstable isotopes?

12. What do all of the isotopes of oxygen have in common in terms of the particles they are made of?

13. What type of particle varies in number in the isotopes of oxygen?

14. Do the number of protons and number of neutrons need to be the same in an atom of a stable isotope?

15. Do the number of protons and number of electrons need to be the same in a (neutral) atom?

## Part 2: Mix Isotopes

1. What are the factors that affect the average atomic mass of a mixture of isotopes?

2. Beryllium (Be) and Fluorine (F) have only one stable isotope. Use the simulation and the Periodic Table to complete the following data table:

Element | Mass of 1 atom | Average mass of 2 atoms (sim) | Average mass of 3 atoms (sim) | Atomic mass (periodic table) |

Beryllium (Be) | 9.01218 amu | |||

Fluorine (F) | 18.99840 amu |

3. Why are all the values in each row of the table above the same?

4. Lithium has only two stable isotopes. Use the simulation to determine the following:

a. Mass of a lithium-6 atom = __________________amu

b. Mass of a lithium-7 atom = __________________amu

c. The average atomic mass of a sample containing ** three** lithium-6 atoms and

**lithium-7 atoms =**

__two________________amu

*(Note: This mixture does not match the mix found naturally for lithium. We will get to the real mix later!)*

d. Is the average atomic mass you just determined closer to the mass of lithium-6 or lithium-7? Explain.

5. Describe a method to calculate the average atomic mass of the sample in the previous question using only the atomic masses of lithium-6 and lithium-7 without using the simulation. (Hint: There are three lithium-6 atoms and two lithium-7 atoms for a total of 5 atoms in the sample.)

6. Test your method by creating a few sample mixtures of isotopes with the simulation and see if your method correctly predicts the average atomic mass of that sample from only the atomic masses of the isotopes and the quantity of each isotope. Use the table below to track your progress. The first row of the table is started with the some of the information for lithium as an example.

Element | Isotopes | Mass (amu) | Quantity in Mixture (number of atoms) | Your Calculation: Average atomic mass of sample (amu) | From Simulation: Average atomic mass of sample (amu) |

Li | lithium-6 | 6.01512 | 3 | 6.41548 | 6.41548 |

lithium-7 | 7.01600 | 2 | |||

## Part 3: Nature’s Mix of Isotopes

Under the “Mixtures” tab of the simulation, click on the “Nature’s Mix” button and then examine the results for several different elements.

1. If you assumed 100 total atoms in a sample, how could you relate the % values shown in the simulation into a number you could use for your calculation of average atomic mass?

2. Calculate the average atomic mass of each of the following elements using your method from above. Test your answer using the Nature’s Mix of Isotopes in the simulation (sim) and the atomic mass of the element listed on the Periodic Table.

Isotope 1 | Isotope 2 | Isotope 3 | Check Answer with Sim | |||||

Element | Mass (amu) | Percent (%) | Mass (amu) | Percent (%) | Mass (amu) | Percent (%) | Calculated average atomic mass
| |

Hydrogen | 1.007 | 99.98 | 2.01410 | 0.011 | – | – | ||

Silicon | 27.97 | 92.22 | 28.9764 | 4.685 | 29.97377 | 3.092 | ||

Nitrogen | 14.00 | 99.63 | 15.0001 | 0.364 | – | – | ||

Argon | 35.96 | 0.336 | 37.9627 | 0.063 | 39.96238 | 99.60 |

Show your calculation for the atomic mass of silicon:

## Part 4: Exercise

1. The atomic mass of boron is 10.81 amu. Boron has two isotopes: Boron-10 has a mass of 10.01 amu. Boron-11 has a mass of 11.01 amu. What is the percentage of each isotope in boron?

a. The two percent abundances need to add up to a total of 100%. Write an equation using variables that shows this relationship.

b. Write another equation that shows the setup for the average atomic mass using your calculation method. Use variables for the unknown, missing information.

c. You now have a system of two equations that you can solve with your algebra math skills! Solve for one of the variables in part a in terms of the other variable, then plug it into your equation in part b and solve for the variable. Plug that result back into the equation in part a to solve for the other variable. Check your answers with the simulation.