Determination of Calcium Carbonate in Eggshells

Introduction
The major component of eggshells is calcium carbonate (CaCO₃). This analysis will be done volumetrically by using a characteristic reaction of carbonate compounds, namely their reaction with acids. Calcium carbonate (limestone) is very insoluble in pure water but will readily dissolve in acid according to the reaction:

      2HCl (aq) + CaCO₃(s) ----> Ca²⁺ (aq) + CO₂(g) + H₂O + 2Cl^- (aq)

This reaction cannot be used directly to titrate the CaCO₃ because it is very slow when the reaction is close to the endpoint. Instead the determination is achieved by adding an excess of acid to dissolve all of the CaCO₃ and then titrating the remaining excess H₃O⁺ with NaOH solution to determine the amount of acid that did not react with the calcium carbonate. The difference between amount of the acid (HCl) initially added and the amount left over after the reaction is equal to the amount used by the CaCO₃. The reaction used to determine the leftover acid is

      HCl (aq) + NaOH (aq) ----> H₂O + Na⁺(aq) + Cl^- (aq)

Stockroom: Things to borrow and return on the same day.

• Buret and brush (~$75)

Procedure

1. Each group of two students should obtain one egg and the necessary glassware from the stockroom.
2. Break the egg into a beaker. Add water to the egg and stir before pouring down the drain.
3. Wash the shell with deionized water and peel off all of the membranes from the inside of the shell.  If you leave the membranes, the proteins in the mebrane will react with the sodium hydroxide.  Dry your eggshell with a paper towel and put into a labeled beaker.  Wash your hands with soap and water.
4. Dry the shell for about 10 minutes in the oven.
5. Grind the shell to a fine powder in a mortar.

• The rest of this experiment is done individually.

1. Accurately weigh between 0.450 and 0.550 g of dried shell into each of 3 labeled 250ml conical flasks. Be certain you record the mass of shell for each flask in your notebook.
2. Add several drops of ethanol to each flask. This acts as a wetting agent and helps the HCl dissolve the CaC0₃.
3. Pour  about 40 mL of 1.0M HCl into a beaker.  Using a syringe pull up 10.0 mL of the 1.0M HCl and then slowly add the 10.0 mL  to each 250mL flask with eggshell in it. Swirl the flask to wet all of the solid. Any HCl that you do not use should be disposed of in the sink and diluted with plenty of water.
4. Heat the solutions in the 250 mL flasks until they begin to boil and then allow them to cool.  Do not let them boil dry! Rinse the walls of the flask with water from your wash bottle.
5. Add 3-4 drops of phenolphthalein (a.k.a PhPL) to each flask.
6. Fill a beaker with about 100 mL of 0.1M sodium hydroxide.  Using a funnel, partly fill a clean buret with 0.100 M NaOH solution to rinse it. Empty the buret into the sink. Fill the buret with the NaOH solution. Run some solution out to remove all bubbles from the tip. Replenish the solution in the buret if necessary. Read and record the initial volume to +/- 0.01 ml.
7. Titrate one sample to the first persistent pink color. When you are close to the endpoint the color will fade slowly. Add the remaining NaOH dropwise (or by half drops) until the color remains for at least 30 sec. Read and record the final volume to + 0.01 ml. 
8. Repeat the titration for the other two samples.
9. Calculate the percent calcium carbonate in each sample and the mean value. Calculate the average deviation from the mean.
10. Wash the egg residue out of the conical flask with hot soapy water and a test tube brush.

Data

Record your starting volume of NaOH and the end volume of NaOH used to titrate each sample.  Each number should have two places beyond the decimal.    For example:

Results

Put your data and results in a neat table.  Show one complete calculation for the % calcium carbonate.  Show the calculation for your average and for your average deviation.  

Calculations

1. Calculate the number of moles of HCl added to each shell sample. This is given by the expression:

moles HCl added= (0.01000 L HCl)*(1.00 moles HCl/liter) = 1.00x10^-2 moles HCl
.

2. Calculate the moles of HCl left in each sample after the reaction with CaC0₃.

moles HCl left = (L NaOH used)*(0.100M NaOH)*(1 mole HCl/1 mole NaOH)
.

3. For each sample determine the number of moles of HCl that  reacted with CaCO₃ by taking the difference between the moles of HCl added and the moles of HCl remaining after the reaction is complete.

moles reacted = moles added - moles left

 

4. The moles of CaCO₃ in each sample is calculated by:

moles CaCO₃ = (moles HCl)*(1 mole CaCO₃ / 2 moles HCl)
.

5. Calculate the percent CaCO₃ in each sample by using

% CaCO₃ = [(moles CaCO₃) *(100.09 g CaCO₃/mole CaCO₃)*(100)] / grams of sample
.

6. Calculate the mean value and the average deviation from the mean.

Sample Calculation

Problem:

In an experiment, 0.500 g of eggshell is dissolved in 10.00 mL of 1.00 M HCl. The volume of 0.100 M NaOH required to neutralize the leftover HCl is 29.70 mL. What is the percent CaCO₃ in the eggshell?

 Solution:

Initial moles HCl = V_HCl M_HCl

.
= 0.0100 L x 1.00 mole/L = 1.00x10^-2 mol
.

moles HCl left = V_NaOH M_NaOH

.
= 0.02970 L x 0.100 mole/L = 2.97x10^-3 mol
.

moles CaCO₃ = moles HCl reacted x (1 mol CaCO₃ / 2 mol HCl )

.

moles HCl reacted = 0.0100 - 0.00297 = 0.0070 moles

.

moles CaCO₃ = 0.0070 mols HCl x (1 mol CaCO₃ / 2 mol HCl ) = 0.0035 mols

% CaCO₃ = ( 0.0035 mols x 100.1 g/mol ) / 0.500 g = 70.0%  (report your answer to three significant figures.)

 

Conclusion:

Using complete sentences, in paragraph form, restate the average percent calcium carbonate in the eggshell and the average deviation.  What were the sources of error in this experiment?  Why did you put the eggshells in the oven?  Are eggshells  a good source of calcium?  Why would you want to avoid eating raw eggshells?

lasted edited by K.Stone on 09/27/2004