Controlled Variables:
- The Distilled Water: The distilled water will be the only water that will be used during the entire experiment. The distilled water was prepared and resented by our teacher.
- The calcium carbonate substance: The calcium carbonate substance was constant throughout the entire experiment.
- The Hydrochloric Acid: The hydrochloric acid was one of the main components of this experiment. Therefore this acid was also prepared and presented by the teacher.
- All Equipment: All the same equipments such as the pipette and the filter paper were used throughout the experiment.
Apparatus:
- Unknown white powdered substance
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50cm3 burette
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3 250cm3 Erlenmeyer flask
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200cm3 for hydrochloric acid
- 0.1M of Sodium Hydroxide
- Burette clamp
- Retort stand
- Distilled water
- Goggles
- Filter paper
- A funnel
- Phenolphthalein indicator
- A scale
- Paper towel
- A spoon
- Pipette
- Pipette sucker
- White paper
Method:
- First set up the retort stand and the retort clamp as indicated by the diagram below. Of course, the flask must be prepared in the subsequent steps. However unlike the diagram, place a white piece of paper beneath the flask so that the color change becomes more easily visible.
- After placing the paper towel and scaling the scale, measure out approximately 10.0 grams of the unknown substance and record the exact mass
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Add 200cm3 of 1.00-mol dm-3 hydrochloric acid and stir until the reaction is complete.
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Filter the solution and withdraw 10cm3 using a pipette and make up to 100 cm3 in a volumetric flask.
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Titrate 10cm3 portions against a standard 0.1M NaOH using the phenolphthalein indicator.
- Record both quantitative and qualitative data.
- Repeat the process of titration three times.
Observation (Data Collection):
Measurements:
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Mass of Calcium Carbonate solution:
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Burette: Calcium Carbonate solution made using 250cm3volumetric flask with an uncertainty of ±0.5cm3
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Pipette: 25.0cm3 of 0.100moldm-3 NaOH(aq) ±0.04cm3
Chemical Equations:
Before we delve into the calculations, it is best that we list the necessary chemical equations of this lab. One must note that in this experiment, that there were two chemical reactions. Firstly, the first experiment occurs when the white substance mixed with the HCl. The second chemical reaction took place when the HCl solution meets with the .1M of NaOH
Qualitative Data:
Description of the substance used and produced
Quantitative Data:
Measurement Recorded During the Experiment
*To note, for the calculations below, the supervisor at the time suggested the best data to use would be the average of the HCl used because the numbers found were very consistent.
Calculations
The following steps were taken so to find the % by mass of CaCO3 in the mixture
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The moles of acid presented in 200cm3of 1.00 mol dm-3hydrochloric acid.
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The moles of acid presented in 10cm3of the acid solution titrated.
0.00056molof HCl
- Multiply the volume present in the volumetric flask.
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The total moles acids remaining after the reaction with CaCO3
- Subtract the value in 4 from that obtained in 1 to find moles acid used.
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The stochiometric equation for the reaction of HCl and CaCO3
Hydrochloric acid+Calcium carbonate→Calcium chloride+water+Carbon dioxide
-Therefore the molar ratio is 1 mole of calcium carbonate to 2 moles of hydrochloric acid
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Using this and the value in 5 to calculate the moles of CaCO3 present. Then calculate the mass of CaCO3
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Using the value found in step 7 and the original mass of the mixture, here is the calculation to find percent of the mixture is made up of CaCO3
Uncertainties:
Total Uncertainty=
Conclusion:
This experiment in general was very interesting. During this lab, I was able to learn of how we can identify the amount of a substance by knowing the concentration and volume of another. We were able to do so as we knew the moles and concentration of sodium hydroxide that had enabled us to find the unknown concentration of calcium carbonate.
Because the actual percentage of CaCO3 was not given, we must concur that the accuracy of this lab has been affected and therefore must be assessed by our uncertainty percentage. The percent of uncertainties represents the desultory errors that may have occurred, which either may make the measurements bigger or smaller than the accepted value, due to imprecise measurement. One way to avoid future random errors, better equipment may be needed, or repetition of the measurements. Despite the uncertainties based on lab equipments and other mediums of measurements, one of the biggest possible sources of error lies in the systematic errors of this lab, especially finding the exact endpoint. Though the pink color should indicate when the endpoint is reached, in trial 2 and 3, the solution turned a faint pink color and my partners and I did not know whether that faint pink indicated the endpoint. This could have resulted in a measurement lower than the actual value. To possibly improve this source of error, I believe more time and trials should be allotted to the students. When I was doing my experiment, we were only given a single class time to follow the procedures, write down our notes, and clean up our experiment. If more time was allotted, then higher quality recordings could be made. This would then give us a clearer idea when the endpoint actually is.
Next, though it is unlikely, to improve this lab we may have to rinse the burettes and other flasks prior to performing this lab. In my chemistry classroom, burettes and flasks are arranged so that the clean ones are clearly seen and labeled while the used ones are in the sink. However it could be possible that another student may have placed an unclean burette in the cabinet full of clean equipments. Therefore, for accuracy purposes, all equipment should be washed using soap and tissues.
Reviewing and understanding the errors can significantly improve this experiment. Though my hypothesis of the concentration being 60% calcium carbonate was wrong, this experiment taught me the value and usefulness of the process of titration.
