This pair of experiments uses gravimetric principles to determine basic chemical data. The first experiment finds the molar mass of magnesium, the second finds the formula of a hydrate. Both experiments are written in a format that can be fitted easily to worksheets. They are written at a low level, providing all steps for calculations. Feel free to modify them to utilize more critical thinking skills. With more advanced students, they should be able to “invent” this procedure themselves, if you give them enough background to work with.
Massing Magnesium Metal
PURPOSE – to experimentally determine the molar mass of magnesium
MATERIALS – Goggles, Bunsen burner, matches, ring stand (with ring), clay triangle, crucible with lid, balance, crucible tongs, magnesium strip, desiccator (and desiccant)
REACTION – 2 Mg + O2
PROCEDURE – Wear your goggles at all times. Hot gasses and dust are formed, and can burn or scratch your eyes. Use crucible tongs to handle hot objects. Do not handle any objects with your fingers, as they add oils (and mass) to anything they touch.
1. Set up your ring stand, clay triangle and burner so that the crucible will rest above the hot part of the flame.
2. Heat the empty crucible with the lid on it for five minutes to drive off any water.
3. Using tongs, place the HOT! crucible and lid in the desiccator to cool.
Once cool, weigh and record the mass of the empty crucible and lid.
4. Place a magnesium strip into the crucible. (Handle with tongs or gloves, or paper towel – but not fingers!) Record the mass of the crucible, lid and magnesium combined.
5. Heat the covered crucible and magnesium over the flame. After 15 minutes, use the tongs to raise the lid slightly, letting some fresh oxygen in. Repeat after another 5 minutes, and once more after another 5 minute period. Be careful not to raise the lid much, or some of the MgO will escape as a white smoke.
6. Extinguish the flame and return the HOT! crucible to the desiccator to cool.
Once cool, weigh and record the mass of the crucible, lid and magnesium oxide.
7. Look inside the crucible to observe the product. Compare its appearance to the original magnesium. If any of the magnesium strip appears to be unreacted, repeat the procedure starting from step 5.
DATA – Record your masses and observations.
Unreacted Mg :
A) Mass of Crucible + Lid: ____________________g
B) Mass of Crucible, Lid, + Mg: ____________________g
C) Final Mass of Crucible, Lid + MgO: ____________________g
CALCULATIONS – Remember your significant figures.
1. Find the mass of your magnesium. (B – A) _________grams
2. Find the mass of reacted oxygen. (C – B) _________grams
3. Find the moles of oxygen (mass of oxygen / 32.00g/mole) _________moles
4. Find the moles of Mg (use the balanced equation) _________moles
5. Calculate the molar mass (mass Mg / moles Mg) _________grams/mole
CONCLUSION – Evaluate your data and results.
How does your measured atomic mass compare to the mass listed on the periodic table?
Calculate the percent difference.
Is your value higher or lower?
What possible explanations can you suggest for this error?
Why doesn’t your value have as many significant figures as the one on the periodic table? Suggest how you might perform a more precise analysis.
The Formula of a Hydrate
Objective – to determine the number of water molecules associated with one molecule of copper (II) sulfate experimentally.
Note: Water has mass. It is imperative that all lab equipment be completely dry before any measurements are made, lest a significant error be introduced.
Caution! The evaporating dish will not look any different when HOT. Use the tongs to handle it from the first moment you heat it until you are finished and have made sure it is cool.
Goggles & apron
Desiccator (with desiccant)
Fine copper (II) sulfate crystals
1. Using the tongs, heat the evaporating dish over the flame to drive off excess water. Cool the dish in the desiccator.
2. Weigh and record the mass of the empty evaporating dish.
3. Place about one spatula of the blue (hydrated) copper sulfate crystals in the evaporating dish.
4. Weigh and record the mass of the evaporating dish and crystals together.
5. GENTLY heat the evaporating dish over the flame until all of the crystals are white (no blue remains). If you heat too intensely, the crystals may “pop” and jump from the dish. Heating too long may also oxidize the sample, resulting in an off-white color. The best results are obtained with low heat and patience.
6. Cool the dish in a desiccator, then weigh and record the mass of the dish and the anhydrous copper (II) sulfate.
7. Add a few drops of water and record your observations.
8. Dispose of the copper sulfate as instructed.
(There should be a labeled waste container for this purpose.)
Data: Record your masses and observations:
X. Mass of empty dish: _____________ grams
Y. Mass of dish + blue crystals _____________ grams
Z. Mass of dish + white copper (II) sulfate _____________ grams
A. Mass of hydrated sample (blue): Y – X = _________ grams
B. Mass of anhydrous product (white): Z – X = _________ grams
C. Moles of copper (II) sulfate: B / 159.61 g/mole = ___________ moles
D. Mass of water removed: A – B = ___________ grams
E. Moles of water removed: D / 18.02 g/mole = ___________ moles
F. Ratio of water to copper (II) sulfate: E / C = ___________
G. Formula of the hydrate: CuSO4 ___ H2O
Conclusion: Evaluate your data and results
Determine the average value (rounded to a whole number) obtained by the class.
Look up the actual formula.
Compare your value and the class average to the actual value.
Which was closer to the actual value?
Is it better to have just one measurement or many when determining a value? Why?
What sources of error were present in this lab?
How might you minimize the potential for error?