1. The equation of the reaction that took place is shown below, in which a grey/white precipitate of calcium carbonate was produced.
CaCl2 (aq) + Na2CO3 (aq) ---> Ca (CO3) (s) + 2NaCl (aq)
Calcium + sodium ---> calcium + sodium
Chloride carbonate carbonate chloride
2. The theoretical mass and the experimental mass are going to be found out in order to see how much calcium carbonate should be obtained theoretically and how much was produced in the real reaction.
To find the theoretical mass of calcium carbonate, firstly we have to find the limiting reagent in the reaction.
The mole ratio from the equation is
CaCl2 : Na2CO3
1 : 1
The actual mole ratio of reagents present is
Mass in g - 4 : 6
Molar mass in g mol ^-1 - 110.98 : 105.99
n = mass - 0.03604 : 0.056609 . Molar mass
Having looked at the mole ratio, it is apparent that since calcium chloride has the lowest number of moles present, it is therefore the limiting reagent.
The limiting reagent calcium chloride is therefore used to calculate the theoretical mass of calcium carbonate that can be obtained:
Theoretical yield = number of moles of limiting reagent x mass of calcium . carbonate
= 0.03604 x (40.08 + 12.01 + 16 + 16 + 16)
= 0.03604 x 100.09
= 3.6g
3. Therefore theoretically the mass of the calcium carbonate that can be obtained is 3.6g. The theoretical yield assumes that everything reacts perfectly, and we are able to recover everything 100%. These ideal conditions are rarely present and so we would expect the actual yield to be less than the theoretical yield for this reason.
To calculate the experimental mass, the following calculation is done:
Experimental mass = Mass of filter paper with the precipitate – Mass of filter paper
= 5.1g – 2g = 3.1g
As expected the experimental mass is lower than the theoretical mass.
4. It is not advisable to use sodium carbonate to calculate the amount of product in the reaction. The theoretical yield depends on the limiting reagent and not the other. Here the two reactants are in a molar ratio 1:1 but the actual molar ratio is 0.03604:0.056609. The sodium carbonate is in excess. It is not possible even under ideal conditions for every sodium carbonate to react to form the product. Therefore it is not advisable to use this.
Conclusion
Having looked at the results, it is clearly seen that the actual experimental mass is less than the theoretical mass yield. This is not an unexpected result. According to my results the actual mass of the precipitate produced was 3.1g, while that of the theoretical mass is 3.6g. The percentage yield of this reaction can be calculated by;
Percentage yield = Actual mass X 100 = 3.1 X 100 = 86% . Theoretical mass 3.6
The maximal yield of a chemical reaction would be 100%, a value that is never reached. Yields about 90% are called very good, yields above about 75% are called good, yields below about 60% are called modest, whilst yields below 30% are called poor. This experiment had a border-line very good yield in regards to these literature guidelines.
In practise the theoretical yield based on the balanced chemical equation is never achieved owing to impurities in reagents, side reactions and other sources of experimental error.
The possible sources of error in this experiment may include:
- Material used may have been tampered with and so would affect the overall results.
- Wrong measurements were taken.
- Error arrising from human judgement.
- The balance only recorded 2decimal points.
- The filter paper may not have been left long enough to dry.
A possible modification to this experiment would be to make the sodium carbonate the limiting reageant rather then the calcium chloride as it was in this case. This would be done so that we would have a smaller number of moles of sodium carbonate then calcium chloride.
Although my experiment was successful, many improvement could have been made to both my experiment and too the experiment. This includes:
- Repeating the measurements for more trials so that more accurate answers could be found.
- Using an accurate method to measure the mass, so as to reduce the errors in the experiment.
- Make sure that none of the compound is accidentally spilled out.
- Use larger quantities so to reduce the error in their recording
