Determination of Empirical Formula of Magnesium Oxide In a compound the atoms of different elements are present in numbers whose ratio is usually an integer or a simple fraction. The simplest (or empirical) formula of the compound expresses that ratio. Simplest formulas are determined by establishing the mass of each element present in a sample of the compound. From those masses one finds the number of moles of each element present. The mole ratio is also the atom ratio in the compound and that ratio provides the subscripts for the simple formula. To find the mass of each element in a compound one must carry out at least one chemical reaction. Sometimes it is possible to form the compound directly from its elements. This is called a synthesis. In this experiment we will synthesize magnesium oxide by heating magnesium in the oxygen in the air: magnesium + oxygen  magnesium oxide By weighing the magnesium before the reaction occurs, and the magnesium oxide produced after, one can calculate the mass of the oxygen that reacted with the magnesium. To obtain good results in this experiment you must make each weighing precisely. Pre Lab Assignment 1. A 0.87 g. sample of silver reacted completely with sulfur and formed 1.00 g of silver sulfide. Find the simplest formula silver sulfide. 2. A 5.00 g sample of aluminum metal is burned in an oxygen atmosphere to provide 9.45 g of aluminum oxide. Determine the simplest formula of aluminum oxide. 3. Find the MSDS for solid Magnesium. List the major health and safety concerns regarding Mg. Is it safe to handle Mg at room temperature? 4. The product that we are trying to produce in this experiment is magnesium oxide. (A) What are the 2 major components of air? (B) Magnesium will react with both components in air to make simple binary ionic compounds of Mg. Write the balanced reactions for both. Purpose 1. To determine the simplest formula of magnesium oxide. 2. To find the percent magnesium in magnesium oxide. Procedure 1. Place a clean crucible and cover on a clay triangle on an iron ring (Fig. 5- 2). The crucible cover should be tilted leaving a small opening. Heat the crucible strongly for about 1 minute to drive off any moisture. Allow the crucible and cover to cool to the touch and then weigh them together. 2. Obtain a piece of magnesium ribbon approximately 50 cm long. Coil the magnesium and add it to the crucible. Weigh the crucible, cover and magnesium. 3. With the cover off (Fig 5-1), heat the crucible. Increase the temperature gradually. When the magnesium ribbon glows red, or ignites, cover the crucible quickly and reduce the amount of heat applied. To prevent any loss of product, the crucible must be covered when you observe ignition. After about a minute remove the cover (to let in more oxygen) and heat until you observe the magnesium glowing or igniting. Then replace the cover and reduce the heat. Continue this procedure until no further reaction occurs. Then tilt the cover, and heat strongly for a few minutes. Let the crucible cool.

4. When the crucible is cool, remove the cover. Use a stirring rod to grind the contents of the crucible into small particles. Rinse the particles remaining on the stirring rod into the crucible with about ten drops of distilled water. Replace the cover, leaving a small opening. Heat gently until the water begins to boil. The water is added to convert any magnesium nitride that might have formed during the reaction to magnesium oxide. Remove the burner, and waft the vapor to see if it has any odor. Then, continue heating until the residue is thoroughly dry. Cool and then weigh the crucible, it's lid and the product, magnesium oxide.

Data Write the observations of combustion of magnesium in your data section. 1. Weight of the crucible and cover

=_________ g

2. Weight of the crucible, cover and magnesium = _________ g 3. Weight of the magnesium = ___________ g 4. Weight of the crucible cover and oxide 5. Weight of the magnesium oxide = 6. Weight of the oxygen

=

= _______ g

__________ g

__________ g

7. Number of moles of magnesium = Mass of Mg x(1 mol Mg/24.31 g Mg)

__________ mol

8. Number of moles of oxygen = __________ mol Mass of O x (1 mol O/16.00 g O) 9. Ratio of moles of magnesium to moles of oxygen Moles of Magnesium : Moles of Oxygen = _______ : 10. Empirical Formula of magnesium oxide = _________

______

11. What should be the formula of magnesium oxide? _____________

Experimental Weight of Magnesium 12. Mass % of Magnesium = ------------------------------------------- x 100 Total weight of Magnesium oxide Actual Weight of Magnesium 13. Mass % of Magnesium = --------------------------------------------- x 100 Total weight of Magnesium oxide Experimental value - Actual value 14. % Error = ---------------------------------------------- x 100 Actual value 15. Calculate the class experimental (precision) error for the percent magnesium in magnesium oxide. Discussion Questions: 1. Based on your experimental data, write the empirical formula for magnesium oxide. 2. Based on their positions in the periodic table: a. What is the most likely charge of a magnesium ion? b. What is the most likely charge of an oxygen ion?

3. Based on the charges of magnesium and oxygen ions (see #1), what would one predict for the formula of magnesium oxide? 4. Is the formula you derived in Question #3 the same as the empirical formula you determined experimentally? If not, explain what might have caused this discrepancy. Was the number of moles of oxygen that combined with one mole of magnesium that you obtained in your experiment larger (ratio less than 1:1) or smaller (ratio more than 1:1) than the accepted value? 5. Based on your experimental data, what is the mass % of magnesium in magnesium oxide. How does it compare with the actual value? 6. Use the formula of magnesium oxide you derived in Question #2 to write out a balanced chemical equation of the burning of magnesium metal in oxygen gas to generate magnesium oxide. Make sure to indicate the physical state (s, l, g, or aq) for each of the substances in the equation. 7. Why is water added to the crucible during the formation of magnesium oxide? 8. Compare the mass of the Mg ribbon with the mass of the magnesium oxide. Notice that the mass of the magnesium oxide is greater than the mass of the Mg. How do you account for this apparent increase in mass?

9. What are primary sources of error/deviation in the experiment? How would factors such as (i) incomplete conversion of Mg3N2 to MgO or (ii) residual Mg(OH)2 in the product affect your results? Does this method appear to be a valid way to determine the formula of metal oxides? 10. Suppose all of the magnesium metal did not combust (burn) to form the product, magnesium oxide. How would that impact (too high or too low) the following? a. Mass of product obtained b. Calculated mass of oxygen used c. Ratio of magnesium to oxygen 11. Suppose you forgot to put the cover on the crucible during the heating and therefore the smoke escaped or you accidentally spilled some of the product from the crucible. How would that impact (too high or too low) the following? a. Mass of product obtained b. Calculated mass of oxygen used c. Ratio of magnesium to oxygen 12. The magnesium was burned in air. We assumed it mixed solely with oxygen, but that is not entirely correct. At high temperatures, magnesium will also combine with the nitrogen in the air to form magnesium nitride (a green-yellow solid). Suppose some of your product massed was magnesium nitride instead of magnesium oxide. How would that impact (too high or too low) the following? a. Mass of product obtained b. Calculated mass of oxygen used c. Ratio of magnesium to oxygen 13. A student did not allow adequate time to evaporate all the water added. What impact would this error have on the final calculations? 14. Determine how each of the following errors would also affect the mass of magnesium oxide at the end of the experiment and how each would affect the ratio of Mg to O. a. Your sample splatters and some of the white powder jumps out of the crucible.

b. You didn’t add enough water to fully react with the magnesium nitride. (Think carefully about this one!) c. If the surface of the Mg ribbon you used were covered with a thin oxide coating prior to the reaction, would your mass percent calculation of magnesium be too high or too low? Explain. 15. Which of the errors discussed were most likely to occur in your trial? Explain. 16. What improvements could be made to this experiment? 17. What are the cautionary measures that you should take in handling the crucible in the experiment? 18. What type of balance did you use, for mass measurements? How many significant figures were in your measurement? 19. What cautionary measures do you take in handling the balance? 20. Using the Merck Index, list two practical uses for magnesium.

21. A 3.70 g sample of sodium is allowed to react completely with sulfur to form a sulfide which weighs 6.30 g. Calculate the following: a) b) c) d) d)

mass of sulfur __________ moles of sulfur __________ moles of sodium __________ ratio of sodium : sulfur ______: _______ empirical formula of sodium sulfide ___________

22. Explain two experimental techniques and/or safety issues that you will apply during your next laboratory assignment. 23. What evidence do you have that a chemical reaction took place? 24. Suggest a modification to the procedure that would be more likely to ensure that all the Mg would react completely with O2. 25. Suggest any extensions for this lab. Think about how you could apply the knowledge that you earned. 26. What two chemistry themes might be incorporated in this laboratory investigation? Explain why they are applicable.