Soda Ash Lab

The goal of this experiment is to have students obtain Potassium and Sodium Carbonates from wood ash. Wood ash contains a mixture of various minerals and compounds that are combined in a heterogeneous mixture. This mixture of minerals contains the wanted Potassium and Sodium Carbonates, and the students will conduct various purification processes in order to extract the Potassium and Sodium Carbonates. This lab experiment also demonstrates the difficulties our ancestors had when extracting these Carbonates. The principle behind how the carbonates can be purified is due to the minerals various solubilities in warm and cold conditions. The process of purification that the students completed was a series of mixture heating and filtrations. Each

The purpose of this experiment is to determine an unknown concentration of acid (hydrochloric acid) with a standard solution of a base (sodium carbonate) using titration method.

The mixture of sodium bicarbonate and sodium carbonate can be separated due to the different temperatures at which the carbonate and bicarbonate decompose. This separation is the focus of the experiment. The bicarbonate decomposes when heated by the Bunsen burner. The carbonate also will decompose but at a much higher temperature. So, when the bicarbonate is heated it decomposes into carbonate, and during the decomposition it releases water and carbon dioxide. Furthermore, the amount of water and carbon dioxide released can be found by finding the difference in mass of the original mixture (before heating) and the final mixture (after heating). The mass of the water and carbon dioxide can be used to find the percent composition of bicarbonate in the mixture.

The purpose of this experiment was to separate the component of three mixtures sand, sodium chloride and calcium carbonate then calculate the percentage by mass of each component recovered from the mixture. The other purpose of this experiment was to show us the students the concepts associated with physical and chemical properties of substances.

In this experiment, the precision of percent by mass of sodium carbonate was decent. It seemed to be consistent, although we seemed to have an outlier in our fifth trial. I believe this was due to human error of adding too much vinegar to this graduated cylinder. The accuracy of our results was decent in comparison to the rest of the class’s data, but our results were on the higher end compared to the averages of the class data, though not too high to be considered

There were three different runs in order to have comparable data and to increase the validity of the experiment. The first run gave 39% of Sodium Bicarbonate. The second run gave 34% of Sodium Bicarbonate while the third one gave 39%, which is the same as the first run. The average percentage was 37% which is much lower than the manufacturer's percentage of 59%. The percent error came out to be -37%.

However, the data found in the experiment did not match the theoretical results at all. It was found in 2 trials the amounts of CO2 produced reached into the negative numbers. Have a negative number produced in a reaction is obviously unfeasible. Many sources of error could have lead to the inaccuracy of the trials such as, splattering of water out of cup when tablet was droped in or when reaction was taking place, evaporation of the water, not all CO2 escaped for the cup, the amount of solution not measured

Throughout the course of the experiment, the weight of the beaker and liquid, the weight of the Alka-Seltzer tablet, the weight of the beaker with liquid plus the weight of the tablet, and the weight of the beaker with all of the contents after the bubbling ceased remained roughly constant and did not vary widely. However, a trend is able to be seen in Figure 1. It is clear that as the mL of vinegar used in each experiment run increased, the mass percent of NaHCO3 increased as well. During the construction of Figure 1, experiment runs four and six were deleted to create the expected graph which consists of a gradual increase and eventually leveling off into a plateau.

After the amount of carbon dioxide lost was found, stoichiometry lets the loss of carbon dioxide to help find the amount of sodium bicarbonate reacted. Moreover, the mass of sodium bicarbonate and the mass of the original tablet help calculate the percent by mass of sodium bicarbonate. An observation made is that the changing amounts in each trial according to Table 1.1 show that, with changing amounts of starting material, the mass of carbon dioxide and the mass of sodium bicarbonate are inevitably affected. However, looking closely at Figure 1.1, the amounts of vinegar added increased in the experiment consequently making it the excess reactant to the sodium bicarbonate causing the graph in Figure 1.1 to level out. The tables should then show a steady increase in the percent by mass in the tablet and the amount of carbon dioxide

My hypothesis is that if I lose the weight of CO2 in Alka Seltzer then I will be able to determine the percent of sodium bicarbonate in the Alka Seltzer. This is because

ii. The second part of the titration series involves titration of NaOH with Hydrochloric acid (HCL). Again, three reps of titration and a blank titration have to be completed. A volumetric pipet is used to measure 10.00mL of HCL into three labeled conical flasks. Then the flasks are filled with deionized water until about the 50mL mark. A buret is

During a titration the pH of the solution will be monitored using a pH meter from that we get a titration curve. The titration curve is then used to determine the equivalent molecular weight and Ka value of the unknown weak acid, from that we are

This showed that there was a limiting reactant in the Alka-Seltzer tablet that forced the production of product to end. When the volume of vinegar reached 10 mL, the average mass of NaHCO3 plateaued equating to an average of 45% by mass of NaHCO3. Due to the high amount of trials, the accuracy in the determination of percent by mass of NaHCO3 increased. However, the standard deviation showed that the precision was low. This brings The percent by mass of NaHCO3 in each tablet in this experiment increased as the amount of vinegar used in solution increased.

The amount of soda ash needed for the experiment was calculated using the following equation: sample weight of unknown=0.1103M (18ml×150.99)/(10×2× %〖Na〗_2 〖CO〗_3 ) An analytical balance was used to weight the calculated amount of soda ash. A piece of weighing paper instead of a weighing boat was used. The mass was recorded. The weighed soda ash was transferred into a 250 mL beaker, then the sample was dissolved in approximately 70 mL of water. The pH meter and electrode was obtained, rinsed with DI water, and calibrated using pH 7 and pH 4 buffer. A burette was obtained, mounted on a ring stand, and filled with the standardized HCl solution, that was prepared in Experiment 2. Since magnetic stirring bars and stirring plates were not available, the students

For this experiment, a pH meter was used so this part of the experiment began with the calibration of the pH meter with specified buffers. The buret was then filled with the standard HCl solution and a set-up for titration was prepared. 200g of the carbonate-bicarbonate solid sample was weighed and dissolved in 100 mL of distilled water. The sample solution was then transferred into a 250-ml volumetric flask and was diluted to the 250-mL mark. The flask was inverted several times for uniform mixing. A 50-mL aliquot of the sample solution was measured and placed unto a beaker. 3 drops of the phenolphthalein indicator was added to the solution in the beaker. The electrode of the pH meter was then immersed in the beaker and the solution containing the carbonate-bicarbonate mixture was titrated with the standard HCl solution to the phenolphthalein endpoint. Readings of the pH were taken at an interval of 0.5 mL addition of the titrant. After the first endpoint is obtained, 3 drops of the methyl orange was added to the same solution and was titrated with the standard acid until the formation of an orange-colored solution. Readings of the pH were also taken at 0.5 mL addition of the titrant.