Gravimetric Determination of Calcium as CaC2O4.H2O

 

Introduction

Calcium is the fifth most abundant elements in the earths crust. While it is primarily found in nature in the form of carbonate (such as chalk, limestone, marble, seashells or pearls) it can also occur as the sulfate (gypsum), fluoride (fluorspar), or a combination of fluoride and phosphate (apatite). It also found in water where it contributes to water hardness (we will discuss this in more detail in another lab). Biologically, calcium is of great importance as a major component of bones and teeth. Indeed the recommended daily allowance (RDA) for calcium is 1g for a normal adult.

In basic solution calcium oxalate (CaC2O4.H2O) is only slightly soluble (Ksp=1.9x10-9). The precipitate, however, will dissolve in acid. In this gravimetric method the calcium and oxalate ions are dissolved in acidic solution and the pH slowly raised until precipitation occurs. If this is done slow enough, large crystals can be obtained. In this case the pH is raised by the thermal decomposition of urea.

Procedure

  1. Wash four labeled filter crucibles and dry for approximately 1 hour at 105C.
  2. After heating place the crucibles in a dessicator to cool down for about 30 minutes and accurately weigh.
  3. Dissolve four 0.5g portions of an unknown solid (make a note of the unknown number) in four 400mL beakers with 100mL of distilled water.
  4. Prepare approximately 300mL of 0.1M HCl and add about 75mL of this to each beaker along with 5 drops of methyl red indicator.
  5. Add approximately 25mL of 4% (w/v) ammonium oxalate to each solution while stirring. Wash the stirring rod into the beaker after stirring.
  6. Add about 15g of solid urea to each sample, cover with a watchglass and boil gently for about 30 minutes (in a fume hood) until the indicator turns yellow.
  7. Filter each solution through a weighed crucible using gentle suction.
  8. Add about 3mL of ice cold water to the beaker and use a rubber policeman to remove all of the filtrate to the crucible. Repeat this with small volumes of ice cold water until all the precipitate has been transferred
  9. Use two 10mL portions of ice water to rinse the beaker and pour through the crucible.
  10. Dry the precipitate initially by suction then by heating in an oven at 105C for 1 to 2 hours or until constant weight is obtained.
  11. Keep the precipitate in a dessicator as much as possible as the solid will absorb water from the atmosphere.
  12. Accurately weigh each precipitate in turn the while storing the others in a dessicator.

 

Calculations

From the mass of precipitate obtained in each trial determine the moles of calcium present and the percent w/w of calcium in the original sample. Determine the mean, 95% confidence interval and relative standard deviation from your trials. Comment on the precision of the results.

Questions

  1. Urea, (NH2)2CO, is thermally decomposed in water to raise the pH by the production of hydroxide ions. Write a balanced equation for the decomposition.
  2. An alternative method for this experiment increases the pH by simply adding the ammonium hydroxide solution. Why might the urea method be prefered?
  3. When 1.00g of the dry calcium oxalate was heated at 210C it was found to decrease in mass to 0.88g. Give an explanation for this observation.
  4. Over what pH range does methyl red change color?
  5. An alternative method for analysis of calcium carbonate is acid-base back titration. Four different samples were analyzed by the two methods and the following results obtained:

Sample number

Gravimetric method (% w/w)

Back titration method (% w/w)

1

55.6

56.2

2

60.4

59.9

3

58.7

56.0

4

55.6

55.3

Do these two methods vary significantly based on these results?

 

 

Reference:

Harris, D. C., Quantitative Chemical Analysis, 3rd Edition, W.H. Freeman and Company.