Name
Date
Pd
Chemistry – Unit 5 Worksheet 1 - Relative Mass
Table for #1 & #2: The density of various gases at standard temperature and pressure:
Gas:
Density (g/L) Relative Mass
Formula
Hydrogen
H2
0.0899
Nitrogen
N2
1.2506
Oxygen
O2
1.4290
Fluorine
F2
1.695
Chlorine
Cl2
2.994
Use the data table on the previous page to answer question 2. Be sure to label your quantities and show all reasoning!
1. a) Determine the relative masses of the diatomic gases in Table 1, relative to the lightest (remember, our data
for combining volumes showed these gases are diatomic molecules in the elemental state). Show your
reasoning below:
b) How do the relative masses you calculated relate to the numbers in the boxes of the elements on the
periodic table? Discuss what you think this means.
Table for #2: The density of various gases at standard temperature and pressure:
Gas:
Formula
Density (g/L) Relative Mass
Hydrogen
H2
0.0899
Helium
He
0.1785
Neon
Ne
0.8999
Argon
Ar
1.7837
Xenon
Xe
5.894
Corrected
Relative Mass
Use the data table above to answer question 2. Be sure to label your quantities and show all reasoning!
2. a) Now let’s examine some other gases. Using the data in Table 1, determine the relative masses of the
monatomic gases relative to hydrogen gas. (hint: monatomic means one atom). Show your reasoning below.
b) How do the relative masses you calculated relate to the mass numbers of the elements on the periodic
table? Looking at the chemical formulas of the gases, do you think the differences can be explained? Use
your explanation to correct the relative masses you calculated for the monatomic gases.
Modeling Chemistry
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You shouldn’t conclude that chemists were able to determine the molar masses of all the elements using this
technique. Measurements of the density of the gaseous phase of many of the elements are difficult, if not
impossible. However, we are going to see that chemists could use another tool – the percent composition of
compounds to determine molar masses.
Part 2: Relative Mass From Compound Data
Many substances combine with oxygen to form a type of compound called an oxide. In the previous unit
we saw that such combinations often occur in multiple proportions. John Dalton made the assumption
that the lowest ratio was a 1:1 combination of elements. For now, we will make a similar assumption.
We may have to re-examine this assumption later.
Based on % composition data, the mass of each of the following elements that combines with 100 grams
of oxygen was determined:
Element
Assumed
Compound
Formula
HO
Mass of element that
combines with 100g
of oxygen
12.5 g
Carbon
CO
75.0
Nitrogen
NO
87.5
Oxygen
O2
100 g
Iron
FeO
349 g
Mercury
HgO
1250 g
Silver
AgO
1349 g
Hydrogen
Dalton’s relative
mass (Question 4)
Actual
relative mass
(Question 5)
1.0
3. If these elements combine in a 1:1 ratio, then the values in the 2nd column could be used to compare the
masses of these elements. As you did in the relative mass lab, divide these values by the mass of
hydrogen to obtain relative masses of the elements. Record these values in the 3rd column. These are the
values Dalton reported for the masses of these elements.
4. a) Compare the value for the relative mass of oxygen and nitrogen you obtained this way to the value you
calculated in Part 1:
b) Why is Dalton’s assumption that H and O combined in a 1:1 ratio in water not justified?
5. Use the correct ratio of H and O atoms in water to determine the actual relative masses of these elements.
How do these values compare to the molar masses in the Periodic Table? Are there any other elements in
the table that do not combine in a 1:1 ratio? Explain.
Modeling Chemistry
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U5 ws1 v1.4 NOLA