Lab In the first step, a standardized HCl solution is used to titrate the mixtur
ID: 726945 • Letter: L
Question
LabIn the first step, a standardized HCl solution is used to titrate the mixture in the
presence of a pH indicator, phenolphthalein, which has an end point pH of about 8. Once
the solution reaches this pH all of the sodium carbonate, but none of the bicarbonate, will
have reacted with the acid. Therefore, the number of moles of sodium carbonate can be
easily calculated knowing the stoichiometry of the reaction and the volume of acid used
in the titration.
The second step entails continuing the titration of the mixture resulting from step
1. Methyl orange, which has an endpoint at about pH 5, will be used as the pH indicator
to monitor the protonation of bicarbonate. The bicarbonate in the mixture reacts with
hydrochloric acid forming carbonic acid (equation 6) which then decomposes to form
carbon dioxide and water (equation 7). Release of CO2 gas into the atmosphere makes
this reaction irreversible
Question:
4. If the pH indicator used for the titration in step 1 changed color at a lower pH than
phenolphthalein, what effect would it have on the mass and wt % of the sodium
carbonate determined in this experiment? Explain
I am not sure why an indicator would play a role in how much the sodium carbonate weighs or its mass?
Explanation / Answer
The process used to determine the concentration of a solution with very high accuracy is called standardizing a solution. To standardize an unknown solution, you react that solution with another solution whose concentration is already known very accurately. For example, to standardize the hydrochloric acid solution we made up in a preceding lab, we might very carefully measure a known quantity of that solution (called an aliquot) and neutralize that aliquot with a solution of sodium carbonate whose concentration is already known very accurately. Adding a few drops of an indicator, such as phenolphthalein or methyl orange, to the solution provides a visual indication (a color change) when an equivalence point is reached, when just enough of the standard solution has been added to the unknown solution to neutralize it exactly. By determining how much of the sodium carbonate solution is required to neutralize the hydrochloric acid, we can calculate a very accurate value for the concentration of the hydrochloric acid. This procedure is called titration. Titration uses an apparatus called a burette (or buret), which is a very accurately graduated glass cylinder with a stopcock or pinchcock that allows the solution it contains to be delivered in anything from a rapid stream to drop-by-drop. Because titration is a volumetric procedure, the accuracy of the results depends on the concentration of the reagent used to do the titration. For example, if 5.00 mL of 1.0000 M sodium carbonate is required to neutralize a specific amount of the unknown acid, that same amount of acid would be neutralized by 50.00 mL of 0.10000 M sodium carbonate. If our titration apparatus is accurate to 0.1 mL, using the more dilute sodium carbonate reduces our level of error by a factor of 10, because 0.1 mL of 0.10000 M sodium carbonate contains only one tenth as much sodium carbonate as 0.1 mL of 1.0000 M sodium carbonate. For that reason, the most accurate titrations are those performed with a relatively large amount of a relatively dilute standard solution. The obvious question is how to obtain an accurate reference solution. For work that requires extreme accuracy, the best answer is often to buy pre-made standard solutions, which are made to extremely high accuracy (and the more accurate, the more expensive). None of the work done in a home lab requires that level of accuracy, so the easiest and least expensive method is to make up your own standard solutions. (In fact, to illustrate the principles of standardization and titration, you don’t even need a truly accurate reference solution; you can simply pretend that a 1 M solution is actually 1.0000 M and proceed on that basis. Your results won’t be accurate, but the principles and calculations are the same.) When you make up a standard solution, take advantage of the difference between absolute errors and relative errors. For example, if your balance is accurate to 0.01 g, that means any sample you weigh may have an absolute error of as much as 0.01 g. But that absolute error remains the same regardless of the mass of the sample. If you weigh a 1.00 g sample, the absolute error is 1% (0.01g/1.00g·100). If you weigh a 100.00 g sample, the absolute error is 0.01% (0.01g/100.00g·100). Similarly, volumetric errors are absolute regardless of the volume you measure. For example, a 10.00 mL pipette may have an absolute error of 0.05 mL. If you use that pipette to measure 10.00 mL, the relative error is (0.05 mL/10.00 mL·100) or 0.5%. If you measure only 1.00 mL, the relative error is ten times as large, (0.05 mL/1.00 mL·100) or 5.0%. One way to minimize the scale of errors is to use a relatively large amount of solute to make a starting solution and then use serial dilution to make a dilute standard solution to use as the titrant. Serial dilution simply means repeatedly diluting small aliquots of known volume. For example, we might start with a 1.5 M solution of a chemical. We use a pipette to take a 10.00 mL aliquot of that solution and dilute it to 100.0 mL in a volumetric flask to make a 0.15 M solution. We then take a 10.00 mL aliquot of the 0.15 M solution and dilute it again to 100.0 mL, yielding a 0.015 M solution. For example, a reference book tells us that the formula weight of anhydrous sodium carbonate is 105.99 g/mol and its solubility at 20 °C is about 200 g/L. A saturated solution of sodium carbonate is therefore about 1.9 M, because (200 g/L)/(105.99 g/mol) = 1.88+ M. In this lab, we’ll make up a 1.5 M solution of sodium carbonate and then use serial dilution to produce a 0.15 M solution to use as our titrant. We’ll titrate our unknown HCl solution once using phenolphthalein as the indicator, and a second time using methyl orange as the indicator. Why two passes with two separate indicators? As it happens, the neutralization of sodium carbonate by hydrochloric acid is a two-step process. In the first step, one mole of sodium carbonate reacts with one mole of hydrochloric acid to produce one mole of sodium hydrogen carbonate (sodium bicarbonate) and one mole of sodium chloride: Na2CO3(aq) + HCI(aq) ? NaHCO3(aq) + NaCl(aq) In the second step, a second mole of hydrochloric acid reacts with the sodium hydrogen carbonate formed in the first step: NaHCO3(aq) + HCI(aq) ? NaCl(aq) + CO2(g) + H2O(I) Each of these reactions has an equivalence point. The first occurs when the first mole of HCl has reacted with the sodium carbonate to form one mole each of sodium bicarbonate and sodium chloride. The pH at the equivalence point of this reaction happens to correspond very closely to the pH range where phenolphthalein changes color, but is well above the pH range of methyl orange. The second equivalence point occurs when the second mole of HCl has reacted with the sodium bicarbonate to form one mole each of sodium chloride, carbon dioxide, and water. The pH at the equivalence point of this reaction happens to correspond very closely to the pH range where methyl orange changes color, but is well below the pH range of phenolphthalein. Figure 11-2. Methyl orange (left) and phenolphthalein at their equivalence points Since these reactions occur with exactly a 2:1 proportion of hydrochloric acid, the amount of titrant needed to reach the indicated equivalence point with methyl orange is exactly twice as much as the amount of titrant needed to reach the indicated equivalence point with phenolphthalein. In other words, if you use methyl orange, you miss the first equivalence point completely. If you use phenolphthalein, you are misled into believing that the first equivalence point is the final equivalence point. For this reason, the neutralization of sodium carbonate with hydrochloric acid is often used as a (literal) textbook example of the importance of choosing the proper indicator. Note that for titrations with multiple equivalence points, it makes a difference which solution is used for the titrant. For example, if we titrate a solution of sodium carbonate by adding hydrochloric acid, we could observe the first (higher pH) equivalence point by using phenolphthalein as an indicator. When sufficient HCl has been added to convert the sodium carbonate to sodium bicarbonate, the phenolphthalein changes from pink to colorless. We could then add some methyl orange indicator to observe the color change from yellow to red at the second (lower pH) equivalence point, when the sodium bicarbonate is converted to sodium chloride. Conversely, if we titrate a solution of hydrochloric acid with sodium carbonate as the titrant, the acid is always in excess until sufficient sodium carbonate has been added to completely neutralize the acid. In this case, we use phenolphthalein as the indicator, because only one equivalence point exists for this reaction, and it is at the higher pH when sodium carbonate is in slight excess. In this lab, we’ll standardize an approximately 1 M solution of hydrochloric acid by titrating a known volume of the acid with a sodium carbonate solution of known molarity
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