Heat a hydrated salt in a crucible and watch the water of crystallisation drive off as the mass falls to a constant value. Record the mass before and after heating, then work out the moles of water lost per mole of anhydrous salt to find the value of x in the formula M·xH₂O.
Theory — Hydrates and Water of Crystallisation
What is a Hydrate?
A hydrate is an ionic compound that holds a fixed number of water molecules within its crystal structure — the water of crystallisation. The formula is written M·xH₂O, where M is the anhydrous (water-free) compound and x is the whole number of water molecules per formula unit. Heating the hydrate drives off this water, leaving the anhydrous salt behind.
Heating a Hydrate
M·xH₂O →(heat)→ M + x H₂O↑
Mass of water lost = mass of hydrate − mass of anhydrous solid
The water leaves as vapour; the mass that disappears is the water of crystallisation
Finding x
Convert the masses to moles and take the ratio. The number of moles of water per mole of anhydrous compound, rounded to the nearest whole number, is x.
Calculation Steps
moles of anhydrous M = mass of M ÷ molar mass of M
moles of H₂O = mass of water lost ÷ 18.02
x = moles of H₂O ÷ moles of M (round to nearest whole number)
Formula of the hydrate = M·xH₂O
Heat to Constant Mass
Heat, cool, and weigh repeatedly until the mass stops changing — this confirms all water has left.
Mole Ratio
The whole-number ratio of water to anhydrous salt gives x. Small rounding is expected.
Gravimetric Analysis
A measurement based purely on mass changes — no titration or instrument needed.
Quantity
How to get it
Unit
Water lost
hydrate mass − anhydrous mass
g
Moles anhydrous
anhydrous mass ÷ molar mass
mol
Moles water
water lost ÷ 18.02
mol
x
moles water ÷ moles anhydrous
(whole number)
Apparatus
The equipment a real hydrate-formula experiment uses. In the simulation these are modelled for you, but the readings correspond to what each instrument would measure.
Crucible & tongs
Holds the hydrate while it is heated to drive off water.
Analytical balance
Weighs the sample before and after heating to find the water lost.
Bunsen burner
Heats the hydrate strongly to remove all water.
Desiccator
Cools the anhydrous solid without re-absorbing moisture.
Clay triangle & tongs
Supports and handles the hot crucible.
Watch glass
Holds the sample during weighing.
Instructions — Running the Virtual Experiment
The Heating Bench lets you choose a hydrate, weigh it, and heat it until all the water is driven off. The Analysis tab helps you turn your measured masses into the formula. Record every reading in your lab report with screenshots.
Part 1 — Heat the Hydrate (Heating Bench tab)
1
Open Simulation → Heating Bench. Select a hydrate from the dropdown and set the starting sample mass. Record the mass of the hydrate (before heating).
2
Click Heat. Watch the water vapour leave and the balance reading fall. Heating continues to constant mass. Record the mass of the anhydrous solid (after heating).
3
Repeat for at least one more hydrate so you have a second data set. Record the hydrate chosen and both masses each time.
Part 2 — Find the Formula (Analysis tab)
1
Open Analysis. Enter your measured hydrate mass and anhydrous mass. The panel shows the water lost, the moles of each, the mole ratio, and the rounded value of x.
2
Record the water lost, moles of anhydrous salt, moles of water, the mole ratio, and the final formula M·xH₂O. Compare the value of x to the accepted formula for that hydrate.
Simulation — Hydrate Heating Bench
Hydrate Gravimetric Virtual LabHeat the hydrate, then analyse the masses
flame
water vapour leaving
Heating drives water of crystallisation off until the mass is constant.
Set up the crucible
Balance readout
HydrateCuSO₄·xH₂O
Mass before heating9.7 g
Current mass9.7 g
Mass after heating— g
Select a hydrate and click Heat.
Enter your measured masses
#
Compound
Hydrate (g)
Anhydrous (g)
Water (g)
mol anh.
mol H₂O
ratio
x
Enter masses and click "Record this result" to log a row.
Formula calculation
Water lost3.50 g
Molar mass anh.159.61 g/mol
Moles anhydrous0.0388 mol
Moles H₂O0.1942 mol
Mole ratio5.00
x (rounded)5
Formula: CuSO₄·5H₂O
Team Questions
Question 1. A hydrate weighs 9.7 g. After heating, the anhydrous solid weighs 6.2 g. How many grams of water were lost? (Type in g)
Question 2. Using 6.2 g of anhydrous CuSO₄ (molar mass 159.61 g/mol), how many moles of CuSO₄ is that? (Type to 4 decimals — e.g. 0.0388)
Question 3. How many moles of water are in 3.5 g of H₂O (molar mass 18.02 g/mol)? (Type to 4 decimals — e.g. 0.1942)
Question 4. Dividing moles of water by moles of CuSO₄ (0.1942 ÷ 0.0388), what is the mole ratio, rounded to a whole number? (Type a number)
Question 5. What is the formula of the hydrated copper(II) sulfate? (Type it — e.g. CuSO4.5H2O)
Question 6. Why must you heat the hydrate to constant mass (repeat heating until the reading stops changing)? (Answer in a few words)
Question 7 — Challenge. If a student stops heating too early and some water remains, will the calculated value of x be too high or too low? (Type high or low)
Example Lab Report
Sample report demonstrating the expected format and level of detail. Use as a guide for your own submission, and include labelled screenshots of the heating bench and analysis for each hydrate.
To determine the number of water molecules of crystallisation (x) in hydrated copper(II) sulfate, CuSO₄·xH₂O, by heating a measured sample to constant mass and analysing the mass of water lost.
Theory
Heating a hydrate drives off its water of crystallisation, leaving the anhydrous salt. The mass lost equals the mass of water. Converting masses to moles and taking the ratio of moles of water to moles of anhydrous salt gives x.
water lost = hydrate − anhydrous · mol = mass / molar mass · x = mol H₂O / mol CuSO₄
The mole ratio of water to anhydrous copper(II) sulfate came out as 5.00, which rounds to a whole number with essentially no deviation. This indicates five water molecules of crystallisation per formula unit. The accepted formula of hydrated copper(II) sulfate is CuSO₄·5H₂O, so the experimental result agrees with the accepted value (0 % deviation in x). Any error in a real experiment would come from incomplete heating (water remaining) or from spattering loss of solid.
Conclusion
The empirical formula of the hydrated copper(II) sulfate is CuSO₄·5H₂O, determined from the mass of water lost on heating to constant mass.
Practice Questions
Show all work. Use water lost = hydrate − anhydrous, moles = mass / molar mass, x = mol water / mol anhydrous. Molar masses: H₂O 18.02, CuSO₄ 159.61, MgSO₄ 120.37, CaCl₂ 110.98 g/mol.
Question 1
A 5.00 g sample of hydrated magnesium sulfate is heated to 2.44 g of anhydrous MgSO₄. Determine x in MgSO₄·xH₂O.
Explain why the mole ratio is rounded to a whole number when writing the hydrate formula.
Hint: water molecules are present in whole-number amounts per formula unit.
Question 5
A student forgets to cool the crucible in a desiccator and it absorbs moisture before the final weighing. How would this affect the measured mass of the anhydrous solid and the value of x?
Hint: a heavier "anhydrous" mass means less apparent water lost.
Question 6 — Challenge
A 4.00 g hydrate sample gives x = 6.4 instead of a clean whole number. Give two experimental reasons the ratio might not come out as an exact integer.
Hint: incomplete heating, spattering loss, weighing error, or impure sample.
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