Lab #1 Andre Grajalez
Date:August 18,2025
Title:Determining The Molecular Formula of Epsom Salts
Aim:To determine the number of moles of water of crystallization (x) present in one mole of
hydrated magnesium sulfate (Epsom salts, MgSO₄·xH₂O).
Materials/Apparatus: Chemicals:
1)Goggles and Lab Coat 1)Epsom Salts
2)1 Boiling Tube
3)Spatula
4)Bunsen Burner Kit
5)Test Tube Holder
6)Digital Balance
Procedures:
1. A clean, dry boiling tube was weighed using a digital balance, and its mass was
recorded.
2. Approximately 3grams of Epsom salt was added to the tube. The tube and the salt
was weighed and the mass was recorded.
3. The tube was heated gently for one minute and then strongly for five minutes.
4. The tube was left to cool to room temperature and promptly reweighed.
5. procedure four was repeated a further five times until the weight remained constant
and was then recorded.
6. All procedures were repeated for a second time to preform a second trial.
7. The final mass of the tube of the anhydrous residue for each trial was recorded.
Results/Observations:
Measurement Mass(g) Trial 1 Mass(g) Trial 2
Empty tube 23.76g 24.01g
Mass of hydrated epsom salt 3g 3g
Tube and hydrated epsom 26.76g 27.01g
salt
Tube and hydrated epsom 25.84g 26.28g
salt after first heating
Tube and contents after 25.71g 26.08g
second heating
Tube and contents after third 25.49g 25.8g
heating
Tube and contents after 25.25g 25.5g
fourth heating
Tube and contents after fifth 25.24g 25.5g
heating
TABLE SHOWING RESULTS FROM TWO TRIALS OF EPSOM SALT HEATING
Calculations:
Trial 1
● a) Mass of hydrated MgSO₄ used
= (Tube + hydrated salt) − (Empty tube)
= 26.76 g − 23.76 g
= 3.00 g
● b) Mass of water in this hydrated salt
= (Mass of hydrated sample) − (Mass of anhydrous MgSO₄)
= 3.00 g − 1.48 g
= 1.52 g
● c) Mass of anhydrous MgSO₄ in this hydrated salt
= (Tube + anhydrous after constant heating) − (Empty tube)
= 25.24 g − 23.76 g
= 1.48 g
Trial 2
● a) Mass of hydrated MgSO₄ used
= 27.01 g − 24.01 g
= 3.00 g
● b) Mass of water in this hydrated salt
= 3.00 g − 1.49 g
= 1.51 g
● c) Mass of anhydrous MgSO₄ in this hydrated salt
= 25.50 g − 24.01 g
= 1.49 g
6. Calculate the relative formula mass (Mᵣ):
a) Water (H₂O)
H = 1, O = 16
= (2 × 1) + 16
= 18g/mol
Calculate the molar mass of water in one mole of hydrated Magnesium Sulphate:
the total molar mass of water is: 7 moles × 18.02 g/mol = 126.14 g/mol
b) Magnesium sulphate (MgSO₄)
Mg = 24, S = 32, O = 16 × 4 = 64
= 24 + 32 + 64
= 120g/mol
10. find the theoretical Formula of Epsom Salts
Epsom Salts has a theoretical formula of MgSO₄7H₂O
11. From this value deduce the percentage Error
● 1. Moles Calculation:
Moles of anhydrous MgSO₄:
Trial 1: 1.48 g / 120 g/mol = 0.0123 mol
Trial 2: 1.49 g / 120 g/mol = 0.0124 mol
● Moles of water (H₂O):
Trial 1: 1.52 g / 18 g/mol = 0.0844 mol
Trial 2: 1.51 g / 18 g/mol = 0.0839 mol
● 2. Mole Ratio (Water to MgSO₄):
To find the number of water molecules per molecule of MgSO₄, we divide the moles of
water by the moles of MgSO₄:
Trial 1: 0.0844 mol H₂O / 0.0123 mol MgSO₄ ≈ 6.86
Trial 2: 0.0839 mol H₂O / 0.0124 mol MgSO₄ ≈ 6.77
● 3. Percentage Error Calculation:
Percentage Error = (Experimental Value - Theoretical Value/ Theoretical Value) *
100%
We can calculate the percentage error based on the ratio of water molecules found:
Trial 1: (6.86 - 7 / 7) * 100% = (0.14 / 7) * 100% ≈ 2.00%
Trial 2: (6.77 - 7 / 7) * 100% = (0.23 / 7) * 100% ≈ 3.29%
Discussion:
Many ionic solids crystallize with a fixed number of water molecules bound in their lattice;
these are hydrates and the water is called water of crystallization. Heating a hydrate drives off
this water, leaving the anhydrous salt and causing a measurable loss in mass. By comparing
the mass of water lost to the mass of the anhydrous salt and using molar masses, you can find
the stoichiometric number of waters, x, in MgSO₄·xH₂O (Epsom salt). The accepted formula
is MgSO₄·7H₂O.
From my results:
3.00 g of H₂O combine with 1.48 g of anhydrous MgSO₄ in Epsom salts (Trial 1) 3.00 g of
H₂O combine with 1.49 g of anhydrous MgSO₄ in Epsom salts (Trial 2)
0.0652 moles of H₂O combine with 0.0124 moles of anhydrous MgSO₄ in Epsom salts (Trial
1) 0.0657 moles of H₂O combine with 0.0125 moles of anhydrous MgSO₄ in Epsom salts
(Trial 2)
5.26 moles of H₂O combine with 1 mole of anhydrous MgSO₄ (Trial 1) 5.26 moles of H₂O
combine with 1 mole of anhydrous MgSO₄ (Trial 2)
Limitation: A limitation encountered was the fact that not all the Epsom salt could be
transferred into the test tube because of human nature and it would affect the end result
calculations because the values could differ.
Source of Error: Rehydration of Anhydrous Salt
A significant source of error in this experiment was the potential for the anhydrous
magnesium sulfate (the dried salt) to reabsorb water from the atmosphere. After heating, the
anhydrous salt is very hygroscopic, meaning it readily attracts and absorbs moisture.
The experimental method employed to determine the water of hydration of Epsom salt is a
form of gravimetric analysis, specifically based on thermal decomposition or dehydration. It
is founded on the principle of conservation of mass and the ability to selectively remove the
volatile water component from an non-volatile solid through controlled heating.
The procedure is to measure the mass of the hydrated compound prior to and subsequent to
the removal of all of the water of crystallization. The change in mass is directly equal to the
mass of water originally with the salt. This allows the calculation of the molar ratio of the
anhydrous salt to the water and thereby the empirical formula of the hydrate. This method is
elementary to chemistry when characterizing hydrated compounds and finding their
stoichiometric composition. The success of this method is dependent upon complete
dehydration without decomposition of the anhydrous compound and prevention of
re-absorption of water.
Examples of Other Salts Which Can Be Tested in This Manner
This gravimetric dehydration process can also be applied to many other hydrated salts, so
long as their anhydrous forms are heat-stable and will not decompose with the temperatures
required to evaporate the water. Some examples include:
Copper(II) sulfate pentahydrate (CuSO₄·5H₂O): Renowned for its distinctive color change
from blue to white upon dehydration.
Barium chloride dihydrate (BaCl₂·2H₂O): A classic hydrate used in first-year chemistry to
demonstrate water of hydration.
Sodium carbonate decahydrate (Na₂CO₃·10H₂O): Also known as washing soda, which when
dehydrated, can be changed to anhydrous sodium carbonate.
Cobalt(II) chloride hexahydrate (CoCl₂·6H₂O): Often used as a humidity indicator since it is
reversible change of color on hydration/dehydration.
conclusion:The experiment succeeded in determining the water of hydration in Epsom salt
with an empirical formula consistent with magnesium sulfate heptahydrate (MgSO₄·7H₂O).
Gravimetric dehydration technique worked perfectly for the quantitative analysis. Sources of
error, such as rehydration of the anhydrous salt, were identified as potential factors that could
impact the accuracy of the results.
Sources:
Royal Society of Chemistry. (n.d.). Water of hydration. Retrieved from [Insert specific URL
if available, e.g., a page on their education section]
Khan Academy. (n.d.). Determining the formula of hydrated ionic compounds. Retrieved
from
https://www.khanacademy.org/science/chemistry/chemical-reactions-stoichiome/stoichiometr
y/a/determining-the-formula-of-hydrated-ionic-compounds
