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Isotopes Exercise
In this lab you will analyze mass spectrometer data for a series of atoms and molecules. Although atoms have only one atomic number, they may have more than one atomic weight.
Isotopes are atoms that have the same atomic number but different mass numbers. The atomic weight of an element is the weighted average of the exact masses of the naturally occurring isotopes. Each isotopic mass is multiplied by the fractional abundance of the isotope, and each result must be added to the others.
Example:
Atomic weight of B =
(Mass of 10B) x (Abundance of 10B) + (Mass of 11B) x (Abundance of 11B) or
(10.01)x(0.199) + (11.01)x(0.801) = 10.81 Atomic Mass Units (Amu)
Answer the following questions.
1. Suppose the mass spectrum of a hypothetical monatomic element X contains a signal at mass number 14 and another at mass number 16.
a. Sketch the mass spectrum assuming the signal at mass number 14 is three times the height of the signal at 16. (You can make your sketch on paper, take a picture on your phone and paste the picture here)
b. How many isotopes are present? Why?
c. What are the fractional abundances of the isotopes?
2. Consider the mass spectrum of neon, Figure 1.
0
10
20
30
40
50
60
70
80
90
100
20
21
22
Relative Abundance
Mass (amu)
Figure 1. Mass spectrum of neon.
The mass spectrum consists of a signal a major signal at 20 amu, a low abundance signal at 21 amu, and another signal at 22. The heights of these signals are proportional to the number of counts of each mass number and represent the natural abundances of the isotopes. Determine the fractional abundance of each neon isotope from the mass spectrum in Figure 1 and calculate the atomic mass for this element.
|
Mass Number |
Measurement with Units |
Abundance |
Atomic Mass (Calculated) ____________ Atomic Mass (Periodic Table)__________
Procedure
Using a metric ruler, measure the peak height of each isotope and calculate the abundance of each atomic isotope for mercury, molybdenum, and bromine.
Make sure to include units with all measurements and the correct precision of the measuring device.
1. Mercury
0
5
10
15
20
25
30
35
198
199
200
201
202
204
Relative Abundance
Mass (amu)
Figure 2. Mass spectrum of mercury.
|
Mass Number |
Measurement (cm) |
Abundance |
2. Molybdenum
0
5
10
15
20
25
30
92
94
95
96
97
98
100
Relative Abundance
Mass (Amu)
Figure 3. Mass spectrum of Molybdenum.
|
Mass Number |
Measurement (cm) |
Abundance |
3. Bromine, Br2
Consider the fragmentation of diatomic compounds. For example, 158Br2 will give a signal at 79 for 79Br and one at 79 for 79Br. Br has two naturally occurring isotopes, the 79 isotope and the 81 isotope. Determine the origin (formula) of each of the signals for the mass spectrum of molecular bromine. Example, 158Br2 is equal to 79Br79Br.
0
10
20
30
40
50
60
70
80
90
79
81
158
160
162
Relative Abundance
Mass (amu)
Figure 4. Mass Spectrum of Br2.
|
Mass Number |
Measurement (cm) |
Abundance |
Formula |
|
XXXX |
XXXX |
||
|
XXXX |
XXXX |
||
|
XXXX |
XXXX |
Calculations
1. Calculate the atomic masses of mercury, molybdenum, and bromine from your data.
2. Compare your results to the actual atomic weights of these elements found on the periodic table. Calculate the percent difference for each.
Discussion Question
1. Why does the mass spectrum of Br2 contain three signals whose heights are approximately in the ratio 1:2:1? What are the origins of these signals?
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1
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