Hydrates are special ionic salts that trap water withing their crystal lattice. Hydrates have a specific ratio of water molecules to formula unit. Hydrates are named by their ionic salt and the number of water molecules they trap.
For example, sodium sulfate decahydrate can trap 10 water molecules per formula unit. Chemical formulas for hydrates always write the water molecules separately at the end after a dot. For example sodium sulfate decahydrate is written as Na2SO4 ∙ 10H2O
There is a simple procedure for determining the chemical formula of a hydrate. We can take a hydrate and heat it to remove the water. By comparing the weight before and after we can determine the chemical formula of the hydrate.
Example:
The formula of magnesium sulfate hydrate is MgSO4 ᐧ xH2O where x is the unknown number of water molecules. In an experiment 8.86 grams of the hydrate is heated until dry. 4.33g of anhydrous (without water) salt remain.
First we know that in the chemical there was 4.33g of the salt.
Next we will find the amount of water originally present. 8.86g - 4.33g = 4.53g of water.
Now, we will follow all of the same steps for determining an empirical formula. As a reminder the steps are:
Change % to grams (we already have grams)
Divide by molar mass of each element
We also know the formula of the salt and the formula of water. So we will treat the salt as one unit and H2O as another unit.
4.33 ÷ 120.32 (mass of MgSO4 = 24.31+32.01+4(16.0)) = 0.03599 mol MgSO4
4.53 ÷ 18.02 (Mass of H2O = 2(1.01)+16.0 = 0.2514 mol H2O
Divide by the smallest number of moles (it will always be the ionic salt)
0.03599 mol ÷ 0.03599 mol = 1
0.2514 mol ÷ 0.03599 mol ≈ 7
Final formula is MgSO4 ᐧ 7H2O