This is not the complete reaction for the titration because bicarbonate is able to accept one more proton. This reaction produces carbonic acid which decomposes to sodium chloride, water and carbon dioxide. NaHCO3(aq) + HCl(aq) ï‚® NaCl(aq) + H2O(l) + CO2(g) These two reactions equate to a 2 to 1 ratio of HCL to Na2CO3 which is expressed in the equation below. 2HCl + Na2CO3 ï‚® 2NaCl + H2O + CO2
The first two reactions are indicative of the two equivalence points of this reaction. Therefore, two indicators are needed to visualize the end-points. The first indicator is phenolphthalein which will turn from the pink ionic form while in a base, to a colorless form indicating the first end-point in this experiment. At this point exactly one mole of HCl has been added per mole of carbonate. This reaction of phenolphthalein occurs from pH 10 to 8.3 which is within 1 pH of the equivalence point for the carbonate to bicarbonate reaction. The second reaction has an equivalence point at ~ pH 3.7. This is where sodium bicarbonate reacts with a proton to produce an excess amount of CO2 very quickly. To visualize this, an indicator that changes color within 1 pH range of the equivalence point is used. For this experiment bromocresol green (BCG), which changes from blue to green when an end-point is reached within the pH range of 5.5 to 3.8.
BCG will turn from green to yellow at pHs below 3.8 which is well below the end-point of the bicarbonate to CO2 reaction. On a titration curve these two end-points would be where the greatest negative rate change in pH occurs on the graph. Because bicarbonate releases CO2 quickly when approaching the second end-point a premature end-point is visualized in BCG, requiring the mixture to be boiled to release this CO2 and allow for the actual end-point to be visualized. Hypothesis: I propose that by knowing the amount, in weight, of soda ash titrated and the amount, in volume, of standardized HCl used the percent of carbonate in unknown 253 can be determined. By increasing the amount of soda ash used the amount of Na2CO3 will increase requiring the standardized HCl to increase along with it. This dependent increase creates a ratio and subsequently the mass percent of carbonate. The independent variable is the mass of the unrefined soda ash and the dependent variable is the mass of the sodium carbonate. The controls used in the experiment are the indicators used, the temperature of the solution and the pressure.
Experimental: HCl that was previously standardized to 0.12 M was obtained. Soda ash was dried appropriately two hours before that lab started. A clean and dry weighing bottle was used to carry the unknown soda ash. The total amount of carbonate in the unknown soda ash was determined by titration. A buret was rinsed with DI water and then 5 10 mL of HCl solution. The buret was filled with HCl. The weigh-by-difference method was used to obtain 0.2 to 0.25 g of soda ash and quantitatively transferred to a 250 mL Erlenmeyer flask. The soda ash was dissolved in 25 mL of DI water and three drops of phenolphthalein indicator was added. The mixture was titrated until the solution was colorless. Then, three drops of bromocresol green indicator was added to the solution. The titration was continued until the premature endpoint was reached. At this point the solution was boiled and then cooled to room temperature. The titration was continued until the actual end-point was reached.
Qualitative: As the first end-point was reached (1 mole of HCl per 1 mole of carbonate) the solution with phenolphthalein turned from pink to colorless. After the three drops of BCG were added to the Na2CO3 solution it turned blue. As the titration continued the solution moved toward a teal hue to green. When the green solution was boiled the mixture moved back to blue. When the titration continued the solutions color went back to green. Discussion: The soda ash samples were weighed out using the gravimetric method and therefore could be measured to the fourth decimal. Each sample was weighed out right before that trails titration.
This was done because as carbonate will slowly react with water to produce bicarbonate. As each sample was titrated for the first end-point the indicator used, phenolphthalein, moved from the ionic form to the non-ionic form, or from pink to colorless. This is an indication that 1 mole of HCl was added per 1 mole of carbonate to give a product of bicarbonate. BCG indicates an end-point between the pH range of 5.5 and 3.8 while the estimated equivalence the HCl and Na2CO3 titration is 3.7.
This is an area of large error margin. After the titration continued a premature end-point was reached and indicated b BCG turning from blue to green. This premature end-point is the result of the bicarbonate reacting with protons and producing a large amount of CO2 while leaving a significant amount of bicarbonate unreacted. As the actual second end-point was reached, between a pH of 5.5 and 3.8, BCG, the second indicator used, changed from blue to green. This indicated that there were 2 moles of HCl added per 1 mole of carbonate. Meaning the reaction reached completion according to the final theoretical expression presented earlier. The average percent of carbonate was 39.33 with a standard deviation of 0.96. This means that the three trails were within a precise range of each other.
Conclusion: The purpose of this experiment was to determine the percent of carbonate present in unrefined soda ash. It was proposed that because the amount of carbonate depended on the amount of soda ash used and the volume of standardized HCl needed to reach end-point that the percent of carbonate could be determined. A standard deviation of .96 was determined with the average percent of carbonate in the soda ash being 39.33%. Because the standard deviation is less than 1 the average percent of carbonate is considered precise and the hypothesis that the percent of carbonate in the unknown 253 soda ash can be determined and is 39.33% is accepted. Recommendations: To obtain a greater accuracy of molarity a pH meter could have been used instead of a visual indicator.
To do this accurately one would have to record the pH vs. the volume of the solution used in a titration curve. Once the rapid change in pH occurs then you reach the equivalence point. Another way to improve the lab would be to run a blank titration, which is to run the procedure without the primary standard and determine the amount of titrant needed to observe a change with the indicator then subtract this from all replicates. Another recommendation is to use a magnet to stir the Na2CO3 mixture as you are titrating to ensure proper mixing speed is constant. Also the second equivalence point is theoretically pH 3.7 and BCG changes color between 5.5 and 3.8. This is an area of huge error margins and a better indicator that can give a more accurate measure of the end-point would be Methyl Orange. This indicator changes from a basic yellow to an acidic red at a pH range of 3.2 to 4.4.