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Using Titration Techniques

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1. If a strong acid and weak acid have the same concentration of acid, explain why they will have a different pH.

2. Explain why it is important to prime the pipette.

3. When using the titration technique explain why the beakers that contain the acid solutions must be dry before the acid is added, but the volumetric flask and Erlenmeyer flasks may have some deionized water drops in them before they are used.

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https://brainmass.com/chemistry/acids-and-bases/using-titration-techniques-74133

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1. If a strong acid and weak acid have the same concentration of acid, explain why they will have a different pH

pH is measured based on the concentration of H3O+ present in the solution. And the concentration of H3O+ depends on the extent of dissociation of the acid. That is, if you have a 1.0 M HCl, the concentration of H3O+ present in the solution will be 1.0 M because HCl is a strong acid and it dissociates 100 %.

HCl + H2O < H3O+ + Cl-

Whereas for 1.0 M of weak acid, for example HA, it will NOT dissociate 100 %.

HA + H2O  H3O+ + A-

The concentration of H3O+ will not be 1.0 M.

2. Explain why ...

Solution Summary

This solution provides answer to a number of questions about titration techniques.

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Determination of Chloride by Precipitation Titration

DETERMINATION OF CHLORIDE BY PRECIPITATION TITRATION

Abstract: Hydrogen Peroxide (H2O2) solution was titrated with a standardized potassium permanganate (KMnO4) solution and it was found that the unknown solution contained 11.3% (wt/vol%) H2O2. were titrated using silver ions as a titrant. These titrations were first used to standardize a AgNO3 solution, then to find the mass percent of an unknown "A" sample containing chlorine (dichlorofluoroescein was used as indicator at equivalence points). The AgNO3 was standardized to .04990 M, with a standard deviation (SD) of 1*10-4, and an RSD of 0.20%. The unknown had an average chloride by mass percent of 33.89%, an SD of 0.29, an RSD of 0.83%, and a 95% confidence interval (CI) of . ±

Data and Results: All mass measurements were made using a Mettler AE100 or an AG204 analytical scale (scale error is ± .0002). 2.1350 g (± 0.0002 g) of AgNO3 solid dissolved in 250 mL of dH2O in a volumetric flask. The molarity was verified by using the AgNO3 solution to titrate ~1 mmol amounts of NaCl dissolved in 50 mL aliquots of dH2O (with dichlorofluoroescein indicator and dextrin). Equation 1 was used to standardize the AgNO3 solution:

MAgNO3(VAgNo3 added) = MNaCl(50 mL) (eq. 1)

The measured amounts of NaCl, the molarity of the NaCl solutions, the AgNO3 volumes added, and the corresponding AgNO3 molarities are shown in table 1. In table 2 can be seen the average AgNO3 molarity found, the SD, and the RSD.

NaCl
(± .0002 g) NaCl (M) VAgNO3
(± 0.05 mL)
AgNO3(M)
0..0594 0.2031 20.31 0.05000
0.05901 0.2017 20.25 0.04980
0.05840 0.01996 20.00 0.04990
Table 1: NaCl and AgNO3 measurements from the standardization of AgNO3.

Mean Molarity SD RSD (%)
0.04990 0.0001 0.20
Table 2: AgNO3 standardization, statistical analysis of titrations.

To determine the mass percentage of chloride in an unknown stock, labeled unknown "A," ~0.1 g of unknown solid was dissolved in ~50 mL water (and DCF indicator and dextrin) and titrated with the standardized AgNO3 solution. Equation 2 was used to find the number of moles of chloride that reacted. Equation 3 was subsequently used to find the true percentage of the unknown A sample that was in fact Cl-. Note that the SD, RSD, and 95% CI are based on the different [Cl-] percentages that were realized in the experiment.

0.0499 MAgNO3(VAgNo3 added) = molesCl- (eq. 2)

[(molesCl-)(35.5gCl/mol Cl)/(Total g unknown A)](100) = Cl- % (eq. 3)

Table 3 shows the volumes of AgNO3 solution added, as well as the original masses of the unknown solid samples that were recorded. In table 4, the average moles of Cl, the average mass percent of Cl in each sample, as well as the RSD and 95% CI are shown.

VAgNO3
(±0.05 mL)

Mass of Unknown (±.0002 g)
19.40 0.1020
19.61 0.1025
19.90 0.1031
Table 3: Volume AgNO3 used to titrate the given samples of unknown (dissolved in 50 mL water)

Moles Cl (mean) Mean Cl%(g) RSD(%) 95% CI
0.0009710 33.89 0.84 ±0.53
Table 4: Statistical values for unknown sample

Discussion: Although the standardized AgNO3 solution was diluted approximately ~10 mL (as determined by approximation via pipet transfer of the overage amount to a graduated cylinder) over the 250 mL volume that was desired, a 0.499 M concentration was observed. A titrant solution that is too dilute has less titrant delivered during each phase of the titration process and as such we would have expected the total volume delivered during the standardization to be higher (thus leading to attainment of a correspondingly lower AgNO3 concentration). That the concentration is not lower than what we observed is difficult to explain because the concentration of a solution that is
2.1350±0.0002 g AgNO3 dissolved in ~260 mL dH2O should yield a molarity of a ~0.4830. There could be numerous reasons for this, though the primary causation likely stems from a random error. Accordingly, identifying reasons why these lower volumes of AgNO3 solution added would be desired. Transfer of the liquid may not have been efficient throughout the titrations, or perhaps some amounts of water evaporated during the titrations. Drips from the stopcock throughout the titration could have also caused slight volume to be lost. Multiple investigators were present for the interpretation of the buret reading, as such, the chance for random error in this regard is lowered.
Indeed the questions become all the more puzzling when we look at the observed state of the solution following addition of the laboratory dH2O from the sink: the solution was cloudy, implying binding from chloride ions that were already present in the solution. This is paradoxical because this should have lowered the amount of available titrant even further, thus lowering the effective concentration of titrant and raising the expected volume delivery amounts. Yet these sort of results were not observed; as discussed previously the volume delivered was in agreement with a molarity that was originally desired. Note that had the volume of the known NaCl analyte during the standardization been diluted, similar effects on determination of AgNO3 would likely not be seen seeing as the final pool of ions in the analyte solution remains the same despite being more dispersed.
Final results attained from the analysis of the unknown yielded generally acceptable results. Although the SD was 0.28 (not shown above in table), which could imply a very high RSD depending on the mean, the RSD was 0.84% and certainly that is to be considered at least acceptable for the purposes of establishing an approximation for the weight percent of the unknown analyzed, in this case it was unknown "A."
Interpretation of the results with regard to the identification of the unknown was inconclusive though examination of common chloride salts shows that magnesium chloride hexahydrate could potentially be the unknown (table 5). Although this analysis can by no method of judgment be considered exhaustive, it is concluded that a single, pure chloride compound comprising the unknown "A" is unlikely unless it is this hydrated form of the magnesium salt. Looking at the common chloride salts that exist as solids, we see that none of them are near the ~33% by mass chloride that was observed in this experiment except the previously mentioned magnesium salt:

Chloride compound Chloride percent by mass (Approximate)
Ammonium Chloride 67
Calcium Chloride dihydrate 48
Cesium Chloride 21
Ferric Chloride 66
Lithium Chloride 85
Magnesium Chloride 75
Magnesium Chloride hexahydrate 34
Potassium Chloride 48
Sodium Chloride 61
Table 5: Common chloride salts and their chloride percentages by mass.

As these salts are relatively inexpensive to purchase, it is a possibility that the sample could have been a pure solid sample. Conversely, although the unknown's appearance was a white solid, and magnesium chloride salts can be white solids, the probability that the unknown "A" was in fact comprised of other compounds in addition to the chloride compound of interest is likely just as high as the probability that it is a pure sample. It would have been beneficial from a confidence standpoint (with regard to identity of the unknown) to attain the results from all laboratory groups so as to produce a statistical analysis of the groups that analyzed unkown "A."
Note that dextrin was used in the experiment as a reagent to inhibit coagulation of the precipitate crystal that forms from the AgCl binding. This was rather important because otherwise the surface area of the average AgCl crystal would be lower as they tent towards less dispersal. Additionally, the adsorption indicator, in this case dichlorofluoroescein (figure 1) is very important in designing an argentomentric titration of this fajans type.

Figure 1: Dichlorofluoroescein structure
The dye is critical because its sensitivity can play a roll in final results attained for concentration and/or weight percent. The charge of the dye and at what pH is very important as this determines its characteristic nature in solution following formation of the precipitates. In this example we have a crystal that becomes positive past the equivalence point so we needed an anionic dye such as the dichlorofluoroescein that was used. To lower coagulation further it is important to have a low electrolyte concentration in the base solutions (both the titrant and the analyte solutions). Contamination by random charged particles in solution will make the hazards of precipitate coagulation even greater due to the "stickiness" of these ions. In this particular case, and perhaps every experiment executed in this laboratory, this could be problematic seeing as analysis of the dH2O has not been done and thus we do not know the concentrations of heavy metals in solution in addition let alone ions like chloride.

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