Analía Bellizzi – Chemistry Classes

Ronald Reagan Senior High School

Empirical Formula of Magnesium Chloride

Empirical Formula of a Magnesium Chloride

Empirical Formula of a Compound LAB

Purpose of the lab:

Calculate the empirical formula of Magnesium Chloride through an experiment:

  • You will combine a known mass of magnesium (Mg) with hydrochloric acid, HCl (aq), to form a compound containing only the elements Mg and Cl (magnesium chloride).
  • The mass of Cl reacting with the Mg will be found from the difference in the mass of the product and the mass of Mg used.
  • The reaction of magnesium with hydrochloric acid is an example of a single replacement reaction. Can you write the balanced equation for the reaction using the known formula of magnesium chloride?


Safety Note:

Hydrochloric acid is a strong acid that is harmful to the skin and especially to your eyes. Wear your safety glasses or goggles during the entire procedure to protect your eyes, and avoid inhaling vapors of HCl during the drying procedure (use a fume hood if possible). The reaction also produces flammable hydrogen gas (H2), so Bunsen burners should not be used while the reaction is in progress.



    • 150 mL Beaker or Conical Flask
    • Hot plate
    • HCl 6M (3 mL)
    • Heating pad (ceramic pad)
    • crucible tongs
    • Mg ribbon
    • digital balance


1)       Obtain a clean, dry evaporating dish and weigh it to the nearest 0.01 g. Record this value on the report form.

2)       Place a small piece of magnesium ribbon into the evaporating dish and record the mass on the report form. From the difference in masses, record the mass of magnesium used. Note: The mass of magnesium should not exceed 0.15 grams, or the product may be difficult to dry

3)       Measure 3 mL of 6 M HCl* in a 10 mL graduated cylinder (or you may use a pipet) and carefully add the HCl solution to the evaporating dish containing the Mg ribbon (Caution: Vigorous reaction!). Allow the reaction to proceed until the reaction is complete, giving a clear solution with no magnesium particles remaining.

4)       Place the evaporating dish on an electric hot plate and heat to nearly boiling. Avoid excessive heat that can cause dangerous splattering of hot HCl!

5)       Heat the solution until evaporation of the water is complete. The white, solid product that remains is magnesium chloride. It is difficult to tell by appearance when the product is completely dry, so we will use a method called heating to constant weight. When the product appears thoroughly dry, carefully remove the hot evaporating dish from the hot plate (you may use crucible tongs to handle the hot dish), allow it to cool, and record the weight on the report sheet under “first weighing.” Then, place the dish back on the hot plate and heat for 10 additional minutes. Allow the dish to cool and record the weight under “second weighing.” If the second weighing agrees with the first weighing, you may reasonably assume that drying is complete. If the second weighing is less than the first weighing, place the dish back on the hot plate and heat for 10 more minutes and obtain a third weighing. Repeat this process until successive weighings agree to within 0.01 g.

An evaporating dish is much better container than a beaker for this process for two reasons:

    1. The solid product spreads out more in the evaporating dish as the liquid evaporates,  facilitating drying.
    2. In a beaker, the vapors of liquid tend to condense on the walls and run back down, inhibiting drying.

6)       From your data, calculate the moles of Mg and the moles of Cl in the product. From the number of moles of each element, determine the empirical formula of magnesium chloride.

7)       Wash your evaporating dish with water (the product may be washed down the sink) and return your equipment to their proper storage locations before leaving the lab.

* The “M” in “6 M HCl” stands for molarity. This is a standard concentration unit in chemistry, and it means moles of solute per liter of solution, mol/L. One liter of 6 M HCl solution contains 6 moles of HCl.


1) Mass of empty evaporating dish   ————– g
2) Mass of evaporating dish + Magnesium   ————– g
3) Mass of Magnesium (2 – 1 )   ————– g
4) Mass of evaporating dish and Magnesium Chloride

After heating and cooling

First weighing   ————– g
 (if necessary)    Second weighing   ————– g
 (if necessary) Third weighing   ————– g
5) Mass of Magnesium Chloride (4 – 1 )   ————– g
6) Mass of Chlorine in Magnesium Chloride (5 – 3 )   ————– g
7) Moles of Magnesium (show your calculation)   ————- mol
8) Moles of Chlorine (show your calculation)   ————- mol
9) Moles of magnesium divided by the smallest moles # (3 sig fig) —————–
10) Moles of Chlorine divided by the smallest moles # (3 sig fig) —————–
11)  Your experimental empirical formula of Magnesium Chloride

(with whole numbers as subscripts)

12) True known empirical formula of Magnesium Chloride ————– 





element mass Divided by Ar Divide by smallest # Mole ratio (whole #)
Magnesium ………./ 24.3
Chlorine ………./35.5






1.       What safety precautions are cited in this experiment?

2.       Why was an evaporating dish more suitable for this lab procedure, rather than using a beaker?

3.       How would your experimental formula of magnesium chloride “MgClx” have been affected if your product was not dried completely before weighing it? Would “x” be too high or two low?

4.       When 6.25 grams of pure iron are allowed to react with oxygen, a black oxide forms. If the product weighs 8.15 g, what is the empirical formula of the oxide?

5.       A compound of nitrogen and oxygen is 30.46% by mass N and 69.54% by mass O. The molar mass if the compound was determined to be 92 g/mol.What is the empirical formula of the compound?
What is the molecular formula of the compound?

6.       How many moles of copper atoms are in 150 g of copper metal?

7.       How many copper atoms are in this amount of copper?

8.       Write the empirical formula for the following compounds containing 0.0200 mole of Al and 0.0600 mole of Cl.

9.       Write the empirical formula for the following compounds containing 0.0800 mole of Ba, 0.0800 mole of S, and 0.320 mole of O.

10.   When 0.424 g of iron powder is burned in an oxygen atmosphere, 0.606 g of a reddish brown oxide is obtained. Determine the empirical formula of the oxide.



Your lab should include: