1. A mixture of 45.0% NaCl (inert) and 55.0% BaCl₂•2H₂O. If 4.165 g of this mixture is heated until all of the hydrate is decomposed, what mass of solid residue will be left?

BaCl₂•2H₂O (s) -> BaCl₂(s) + 2H₂O(g)
NaCl(s) -> NaCl(s)

We are rounding our answer to 4 significant figures because the least accurate measurements of 5.000 g and 60.00% are 4 significant figures. In the lab manual problem, round to 3 significant figures because the percentages are only 3 sig figs. For the lab manual problem 1, your mixture's mass and percentages are different but the procedure is the same. Note: In a spreadsheet program, the formula at cell K2 would be =A2*B2*F2*H2/D3

2. An ore contains 3.50% ZnS. How many kilograms of zinc are in 880. pounds of the ore? Use molar masses and moles to solve the problem.

A vein of zinc sulfide (mineral name is sphalerite)

Zinc sulfide crystals

Zinc sulfide glows under UV light

Again, the percentage of 3.50% is assumed to be "by weight." 3.50% w/w meaning 3.50% weight to weight is a more explicit way to write it. The main reason you know it is by weight is that the problem is giving the mass of the ore. Also, notice that the pounds "880." has a decimal point. That means it was measured to the nearest pound and the zero is therefore significant. Without the decimal point, "880" is assumed to be measured only to the nearest 10 pounds and only the 88 digits would be significant. Since "3.50%" and "880." both have 3 significant figures, your answer needs to have 3 significant figures.

Practice for 2: A dump truck has a load of ZnS ore that weighs 4.76 tons. An assay of the ore shows a concentration of 5.34% w/w zinc sulfide. How many kilograms of zinc could be extracted from this ore?

3. A sample contains a mixture of magnesium carbonate and sand (an inert material). If 5.00g of sample loses 1.00 g upon heating, calculate the % magnesium carbonate in the sample.

MgCO₃(s) -> MgO(s) + CO₂(g)
Sand(s) -> Sand(s)

Note: The crucible you used for this lab is partially made from MgO (magnesium oxide). MgO most likely came from the heating of magnesium carbonate.

Magnesium carbonate is used for many things. It's used in antacid tablets (TUMS) and in toothpaste. It's put in table salt to keep the salt dry. Gymnasts and rock climbers put it on their hands to prevent slippage. It's also used in plastics and tires. Below is Lake Salba in Turkey. Outcroppings and much of the beaches are made of magnesium carbonate and sand. Mining of it is being considered. To be profitable, the concentration of magnesium carbonate needs to high enough to offset the costs of separating it from the sand. Let's say the concentration needs to be over 35% by weight to be profitable.

Practice for 3. A cup of the sand/magnesium carbonate mixture was scooped up from the beach and dried. The dry mixture weighed 435.6 grams. After heating to 700°C for 25 minutes to convert all magnesium carbonate to magnesium oxide and carbon dioxide, the mixture was reweighed and found to be 310.2 grams. Will it be profitable to mine this mixture to extract out magnesium carbonate (In other words is the % magnesium carbonate over 35%)?

4. Unknowns are prepared using Epsom salt, MgSO₄•7H₂O, and NaCl (inert). Calculate mass percent MgSO₄•7H₂O in a sample from the following data:

MgSO₄•7H₂O(s) -> MgSO₄(s) + 7H₂O(g)
NaCl(s) -> NaCl(s)

Epsom salt is used as a laxative and for creating a bath for soaking feet or body. It is named after the England town of Epsom where people came to drink from a well that had dissolved magnesium sulfate. An unscrupulous supplier may mix the Epsom salt with the cheaper table salt (NaCl) to increases profits. The procedure of weighing a sample, heating it, and then reweighing it can reveal if the Epsom salt was diluted with salt and by how much.

The calculation for row 3 is a simple subtraction.

To find row 7, you will use the 3rd heating value of 53.661 because that is the Epsom salt with the least amount of water. The easiest way to find mass loss is to take row 1 (mass crucible and sample) and subtract row 6 (mass of crucible and residue).

The mass loss is the water coming off. The strategy is if you know the mass of the water coming off, you can turn that into moles of water that was lost. The balanced equation shows us that the number of Epsom salt molecules is 1/7 of the number of water molecules that come off. So dividing the moles of water by 7 gives us the moles of Epsom salt. After you get moles of Epsom salt, you can turn that into grams of Epsom salt. From there it's just division by the starting mass (row 3) to get the fraction of Epsom salt. So this is just like problem 3 above. Below is the spreadsheet layout of the problem. Plug in your values from rows 7 and 3.

5. A pure sample of hydrated cobalt (II) chloride weighing 2.854 g is heated to a constant mass of 1.558 g. Calculate the formula of the hydrated salt.

CoCl₂• nH₂O(s) -> CoCl₂(s) + nH₂O(g)

5) Cobalt chloride is used as moisture indicator.  CoCl₂ is blue and when it gets moist, it turns to red.

Like in the previous problem, the loss of mass is attributed to water vapor escaping as the salt is heated. Like before, if you know mass of water, you can turn that into moles of water. The final mass of 1.558 g comes only from COCl₂, so you can turn that into moles of CoCl₂.

Practice for 5. A white powder is found at a crime scene. Initial analysis shows a powder to be a hydrated form of calcium sulfate, but which form requires the same procedure as used in this lab. The difference between gypsum (drywall) and plaster of Paris (plaster for casts) is the number of water molecules attached to calcium sulfate.
Gypsum is CaSO₄• 2H₂O
Plaster of Paris is CaSO₄• ½H₂O

You start off with a sample that weighs 37.8779 g. After heating until the mass doesn't change, the residue is 29.9512 g. Is it gypsum or plaster of Paris? The answer gives a clue if the suspect is a drywaller (gypsum) or works at a hospital creating casts for broken limbs (plaster of Paris).