CaCO3 in egg shells

A versatile 'back-titration'

This practical can be adapted to determine the amount of calcium carbonate in a variety of different substances ranging from eggs shells to the shells of sea creatures, such as crabs, or in natural or artificial stone such as marble, limestone cliffs or lime mortar from old buildings. The basic reaction is simply:

CaCO3(s) + 2HCl(aq) → CaCl2(aq) + CO2(g) + H2O(l)

It is not practical to measure the exact amount of acid required for a known mass of the sample as it is difficult to determine exactly when all the calcium carbonate has reacted. This practical makes use of a technique known as ‘back titration’ [1]. A known excess of acid is added to the sample ensuring that the reaction goes to completion. The excess acid is then diluted with water, made up to a known volume and then titrated against a standard solution of sodium hydroxide. The amount of acid that reacted with the sample can then be deduced.

Teacher’s notes

This is actually a very versatile practical. Several students each year basically incorporate it into extended essays to answer such research questions as ‘Do free-range eggs contain more calcium carbonate in their shells compared to those produced by battery hens?’ or ‘How does the composition of a particular lime mortar change over time?’ It is a good practical to assess for both Data collection and processing (DCP) and Conclusion and evaluation (CE). There is a reasonable amount of data to collect and although there is no graphing the processing is slightly more complex than a normal titration. The students can put into practice all they have learned about uncertainties associated with their apparatus (balance, pipette, burette and volumetric flask) to arrive at an overall uncertainty. They can then look in the literature to find typical percentage compositions of the material they have used. Perhaps much more important than this is to see whether they can question some of the underlying chemical uncertainties associated with their result during their evaluation. How many students will realise that the result assumes that the only thing in the sample reacting with the hydrochloric acid is calcium carbonate. What would happen if magnesium carbonate is also present (as it is in dolomitic rock) or if other anions such as sulfide (S2-) or sulfate(IV) (sulfite) (SO32-) are also present?

I have used quite small quantities to save on chemicals and the environment. Provided an accurate analytical balance is available to be used this should make little difference to the accuracy of the final result. If you want to use more traditional amounts and 25 cm3 pipettes etc then scale up the amounts. This means weigh out about 1.500 g of the sample and add 50.0 cm3 of the 1.00 mol dm-3 hydrochloric acid. After the reaction, dilute with water and transfer to a 250 cm3 volumetric flask and make up to the mark with distilled water. Then take 25.0 cm3 sample of this for titration with the standard 0.100 mol dm-3 sodium hydroxide solution.

Student worksheet

TO DETERMINE THE PERCENTAGE BY MASS OF CALCIUM CARBONATE IN AN EGG SHELL

INTRODUCTION
This practical is to determine the percentage by mass of calcium carbonate in an egg shell but could in fact be used to determine the amount of calcium carbonate in other samples such as limestone rock or sea shells.

In addition to the stated aim of determining the percentage of calcium carbonate in the shell you are also introduced to the technique of ‘back titration’. Calcium carbonate reacts with hydrochloric acid but it is difficult to determine exactly when all the solid calcium carbonate has reacted to one drop of standard hydrochloric acid solution if it is titrated directly. In this experiment a known amount of excess acid is added to the sample to ensure that all the calcium carbonate has reacted. The excess acid is then diluted and made up to a known volume. Aliquots of this diluted excess acid solution are then titrated with a standard solution of sodium hydroxide.

ENVIRONMENTAL CARE: Because the samples are natural materials and the acids used mainly react to form either calcium chloride or sodium chloride there are no particular environmental issues and the waste can be disposed of down the sink. To save on distilled water this practical uses smaller amounts than more traditional titration experiments.

SAFETY: There are no particular safety hazards except for the usual need for care when handling glassware and 1.0 mol dm-3 strength acids.

ASSESSMENT: This practical will be assessed formally for Data collection and processing (DCP) and Conclusion and evaluation (CE).

PROCEDURE
Carefully wash the shell of an egg to remove any dirt and organic matter attached to it. Dry the shell either in an oven or by using hot air from a hair dryer. Grind the shell into small pieces and weigh accurately about 0.6 g of the shell into a conical flask. Using a graduated pipette add 20.0 cm3 of 1.00 mol dm-3 hydrochloric acid solution. Add the acid slowly and swirl the flask to prevent any small amount of liquid escaping from the flask with the carbon dioxide that is produced. Once the reaction has completely finished add about 20 cm3 of distilled water and transfer all the contents of the flask to a 100 cm3 volumetric flask. Use more distilled water to ensure all the contents have been transferred and to make the final volume to exactly 100 cm3. Take 10.0 cm3 aliquots of this solution and titrate them with standard 0.100 mol dm-3 sodium hydroxide solution using phenolphthalein as an indicator.

Record all your data in an appropriate way and work out the percentage of calcium carbonate by mass in the sample that you took. Evaluate your experiment fully.

calcium carbonate in shells


Footnotes

  • 1. This technique is not specifically stated on the syllabus but all the calculations involved easily fall under assessment statements in Topic 1 : Quantitative chemistry

Comments 5

michael sugiyama jones 17 July 2013 - 13:35

I've tried this practical using 20.0 ml of HCl with 0.6 g of egg shell (without making it up to 100ml with distilled water), and then titrating it with 1.0M NaOH.
What is the advantage of making up the solution to 100 ml?

Geoffrey Neuss 21 July 2013 - 20:10

Hi Michael. I think there are two advantages. The main advantage is that by making it up to 100 cm3 you can then do at least three titrations (one rough and two accurate) to get an average whereas I think the way you are doing it will only lead to one titration? The second advantage is that you will need a larger volume of the 0.100 mol dm-3 NaOH which will give less uncertainty than if you use the smaller volume of 1.00 mol dm-3 NaOH required. I hope this is helpful.

michael sugiyama jones 24 July 2013 - 12:05

Hi Geoff, there is a post about this on the Chem forum of the OCC. Maybe you could check it out and let us know your thoughts?
Thanks,
Mike.

Geoffrey Neuss 24 July 2013 - 20:37

Mike, The sample calculation you have done on the OCC discussion board looks fine. One very small point is that chemists now tend to say 'amount' rather than 'number of moles' (e.g. the amount of HCl in 20.0 cm3 of 1.00 mol dm-3 HCl(aq) = 2.00 x 10-2 mol). It is no big deal but your students should know as that is how it is expressed in the exam papers. Best wishes, Geoff

michael sugiyama jones 25 July 2013 - 04:29

Thanks for the advice Geoff.


To post comments you need to log in. If it is your first time you will need to subscribe.