The purpose of this experiment is to examine the percent yield of a precipitate in a double displacement reaction. A solution of calcium citrate and sodium carbonate were mixed together, then the products were filtered out as so only the precipitate remained. The filtered paper was then dried and the mass of the precipitate in the experiment divided by the theoretical mass of the precipitate from the calculated gave the percent yield. The percent yield that was acquired is about 69.1%.
The chemical reaction that took place was that sodium displaced calcium in a solution of calcium nitrate, and that calcium displaced sodium in a solution of sodium carbonate. The resulting products were solid calcium carbonate and aqueous sodium nitrate (equation 1).
To calculate the theoretical yield, the moles of sodium carbonate and calcium nitrate were calculated to find the limiting reagent. Then the moles of the limiting reagent (calcium nitrate) were used in ratio with the precipitate ( 1:1), calcium carbonate, to find how many moles of the solid were produced, then to conclude the mass of the precipitate (equation 2).
The percent yield was calculated by taking the mass of the precipitate that was observed in the reaction(actual yield), divided by the mass of the precipitate that was hypothesized in the calculations (theoretical yield). This number was multiplied by a hundred to determine the percent yield (equation 4).
Na2CO3 (aq) + Ca(NO3)2 (aq) CaCO3 (s) + 2NaNO3 (aq)
n = OR mole =
Percent Yield = × 100
The law of conservation of mass states that the mass of the reactants must be equal to the mass of the products; no mass can be created or destroyed. In relation to this experiment, the theoretical yield is the calculated mass based on if the result has a percent yield of 100%.
The theoretical yield of the precipitate is mass of approximately 0.695 g (appendix A). Therefore if the percent yield is close to 100% (an actual yield ≈ 0.694), then less experimental errors occurred and the more accurate the result.
1. To ensure safety, goggles were worn.
2. A 150ml beaker with a mass of 68.2g was filled with sodium carbonate until the total mass was 69.08g. The mass of the sodium carbonate contained was calculated and recorded to be 0.88g. 3. Another 150ml beaker with a mass of 72g was filled with calcium nitrate until the total mass was 73.14g. The mass of the calcium nitrate contained was calculated and recorded to be 1.14g . 4. Using a graduated cylinder, 50ml of distilled water was added to each of the beakers
5. A different stirring rod was used to dissolve each of the solids in the beakers. 6. Then the calcium nitrate solution beaker was poured into the other beaker that contained the sodium carbonate solution and mixed thoroughly with a stirring rod. 7. A filter paper, massed to be 0.75g, was subsequently set up (diagram 1) 8. The filter paper was placed inside a funnel and the funnel was placed into a clean 250ml flask.
9. The products of the mixed solutions were slowly and hesitantly poured into the funnel, until nothing remained in the beaker (diagram 2) 10. Then the funnel paper was removed and set to dry
11. The day after, the funnel paper with the precipitate was weighed to be 1.25g and subtracted from the mass of the funnel paper by itself to get an actual yield of 0.48g of precipitate.
Diagram1 : Fold at the dotted line.
Open with three folds on one side and
place in the funnel with a few drop of
water so it sticks to the funnel.
Table 1: The recorded masses of materials and results in the experiment Material
Beaker 1 with sodium carbonate
Beaker 2 with calcium nitrate
Dried precipitate with funnel paper
Calcium carbonate precipitate
The percent yield was calculated to be approximately 69.1%. This was concluded by dividing the mass of the mass observed in the reaction divided by the mass calculated that was theoretically supposed to be produced; actual yield divided by theoretical yield (appendix A).
The percent yield of the calcium carbonate is 69.1%. This means that less precipitate than was expected was produced due to systematic errors that occurred, causing deviation and a result that is inaccurate.
A systematic error that is hypothesized to have occurred is that some of the precipitate made it through the filter. When the two solutions were mixed, a precipitate of calcium carbonate was formed. The precipitate with the products when filtered through the filter funnel must not have had enough time for the reaction to reach its completion. The observation that the flask with the filtered products was foggy only helps to confirm the point that some of the precipitate made it through the filter.
This causing a decrease in the actual yield, consequently leading to a decrease in the percent yield. A solution to this error for future reference would be to use a catalyst to speed up the reaction, or by simply giving the reaction more time so that the net ionic equation of the reactants could form into a precipitate completely .
Another systematic error that may have occurred is loss of material during transfer. In many instances, substances had to moved from one container to another, stored, poured, and all these transfers are imperfect. Substances could have stuck to the sides of the beakers, to the stirring rods, to the tissue paper under the drying filter paper and precipitate.
This excessive loss of matter causes a decrease in the actual yield of the precipitate, and coherently leading to a decrease in the percent yield, lowering the accuracy of the experiment. For this problem to be avoided, using more advanced tools and being more careful when transferring products would be helpful.
A third systematic error may be due to the result of evaporation. When the two solutions of sodium carbonate and calcium nitrate were mixed, it is not ensured that all of the reactants that could be produced (according to the limiting reagent) were reacted. Some of the mass could have been lost and evaporated due the warm temperature in the room due to the season, and the energy produced from the large mass of people in the room and their friction.
This causes a mass of the precipitate to be less than expected and therfore decreasing the percent yield accuracy. To avoid this, covering the reactant beakers would create an equilibrium, causing the amount of reactant that is evaporating equal to the amount of reactant that is condensing back in (figure 1), or by decreasing the temperature from the thermostat in the room so that the room temperature is below the boiling point of the solutions.
Applications to Real World:
Percent yield is essential the chemical manufacturing and engineering business. When the reactants are loaded into the reactor processor, the actual yield is taken and divided by the theoretical yield. In terms of benefits to the company and their economic status, having a higher percent yield would be a plus for the company, saving the company valuable money on reactants and resulting in a higher net profit; the highest amount of products produced , in the least amount of money.
It also would be more environmentally advantageous to them because less fuel and materials are being wasted, which could be used as an advertising strategy as using their products would be environmentally friendly in the eyes of the consumer . Therefore, it would be considered the chemist’s job to improve the percent yield of his or her company, similarly to how this experiment was conducted, and how underlying errors were identified to improve future results.
Calculating Percent Yield
Na2CO3 (aq) + Ca(NO3)2 (aq) CaCO3 (s) + 2NaNO3 (aq)
n (Na2CO3) =
= = 8.303 × 10-3 mol of sodium carbonate
n (Ca[NO3]2) =
= = 6.948 × 10-3 mol of calcium nitrate
moles of CaCO3 =0.0069478
Theoretical yield of CaCO3 = 6.95 × 10-1g
Percent Yield = × 100
Percent Yield of Calcium Carbonate = ×100
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