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Acid/Base Titration Essays

Experiment 2 Acid/Base Titration John J. Purdue CHM 321 – Fall 2012 TA: Scott Cole Section 1 September 4, 2012 Unknown Concentration: X. XX ± X. XX M (@95% confidence interval) (adapted from a report prepared by N. Skrynnikov, 2009) Abstract The concentration of an unknown acid (HA) solution was determined by titration with a standardized solution of sodium hydroxide. The standardization of NaOH was done by titration with a solid acid sample, potassium hydrogen phthalate (KHP), and phenolphthalein indicator. The unknown concentration (Cunknown) was determined to be X.
XX ± X. XX M at a 95% confidence interval, and the methods described herein constitute a simple and reproducible technique that may be applied to quantitatively assess many different acid/base pairs, provided that an appropriate indicator is used to determine an endpoint. Introduction This laboratory exercise relies on a titration  technique to determine an unknown concentration of monoprotic acid in solution. In the process of titration, a basic solution is gradually added to the acidic solution until complete neutralization is obtained.
The ‘end point’ of the titration is detected with the help of an indicator  as color of the solution changes upon neutralization. By measuring the volume of the titrant required to reach the ‘end point’, it is possible to relate the concentration of the acid to the concentration of the base. In this manner, the unknown  concentration can be expressed through the known concentration. The concentration determination is repeated several times in order to improve the precision of the measurements and to estimate the experimental error.
The experiment involves two steps: (i) Standardization of sodium hydroxide (NaOH) solution using potassium hydrogen phtalate (KHP) solution, and (ii) titration of an unknown monoprotic acid solution using the standardized NaOH solution. The two steps, (i) and (ii), are essentially similar. Therefore, only the first step is briefly described below. The neutralization reaction proceeds as follows: NaOH+KHP>Na+ +K+ +P2- +HO Once this reaction is complete, an excess of NaOH starts building up, triggering the response from the indicator: NaOH + HIn(colorless) >  Na+ + In-(pink) + H A question that may arise is why step (i) is needed at all.
Indeed, one could envisage a simpler measurement scheme where the solution of NaOH is prepared with known concentration and used to titrate an unknown acid. Bear in mind, however, that NaOH is a poor primary standard: it is highly hygroscopic, chemically unstable (reacts with CO of air), typically low-purity (if purchased cheap), and has low molecular weight (which leads to higher relative error when the compound is weighed out). Conversely, KHP has many desirable characteristics which make it a good primary standard. This dictates a choice of the two-step scheme, with KHP as a primary standard and NaOH as a secondary standard.
A flow chart of this methodology is shown below in Scheme 1. Scheme 1. General layout of an Acid/Base Titration Experiment Procedure Laboratory procedures were carried out according to the laboratory manual exercise entitled “Experiment 2 – Acid/Base Titration” accessible on the Blackboard Learn course website. All recorded data, including KHP mass and NaOH volumes, are included below in the “Results and Discussion” section of this report. REMEMBER: If any deviations are made from the outlined procedure cited above, they MUST be included here in great detail! Results and Discussion Standardization For the standardization step, the KHP olution was prepared by weighing out 4. 8149 g of (dried) KHP and dissolving it in distilled water to a volume of 250 mL. Considering that the molecular weight of KHP is 204. 23 g/mol, the concentration of the KHP solution (CKHP) is: CALCULATION for CKHP SHOWN HERE! NEATLY AND WITH UNITS SHOWN EXPLICITLY. The titration (standardization) results using 25. 00 mL aliquots of the KHP solution are summarized in Table 1 below. Table 1: Volume data from NaOH standardization measurements Trial 1* Trial 2   Initial Volume (mL) X. XX X. XX Final Volume (mL) XX. XX XX. XX Titrant Volume added (mL) – Vendpoint XX. XX XX.
XX Trial 3 X. XX XX. XX XX. XX * Trial 1 was preceded with the scout titration (trial 0). The results from the scout titration are not included in this table since they are not quantitatively accurate. Already a cursory inspection of Table 1 shows that the results are highly reproducible – the uncertainty in the volume of the titrant is on the order of 0. 1 mL (0. 4%). The data from Table 1 was used to determine the concentration of sodium hydroxide solution, CNaOH . The molar balance conditions corresponding to the complete neutralization (end point of the titration) can be written as: CKHP  [mol/L] VKHP L] = CNaOH [mol/L] Vendpoint [L] The concentration, CNaOH, was calculated on the basis of this formula, using CKHP = X. XX M, VKHP = XX. XX mL, and Vendpoint = XX. XX mL, as listed in Table 1. The values of CiNaOH calculated in this manner were as follows, for each of three trials (i =1, 2, and 3) respectively: X. XX M, X. XX M, X. XX M. The mean concentration is therefore X. XX M. This data/calculation could also easily be organized in a separate table; however, care must be taken to make sure it is neat, labeled clearly, and discussed appropriately. The ncertainty in CNaOH was calculated according to the following formula: , where Here N is the number of measurements, N = 3 , and t is the so-called ‘Student’s coefficient’, t =X. XX assuming that the confidence level is 95% and that the number of degrees of freedom is (N-1) = 2 . The sample standard deviation, s, was calculated according to the above formula using the data from the individual trials, CiNaOH , and their mean. This calculation produced s = X. XX M After rounding off the result appropriately and retaining significant digits, the value of CNaOH obtained was X.
XX M ± X. XX @ 95% CL. This concentration of sodium hydroxide was then carried forward in subsequent calculations to determine the concentration of the unknown. Titration of the Unknown The titration results using standardized NaOH solution are listed below in Table 2. Table 2. Volume data from the titration of unknown monoprotic acid using standardized NaOH solution. Trial 1* Trial 2 Trial 3 Initial Volume (mL) X. XX X. XX X. XX Final Volume (mL) XX. XX XX. XX XX. XX Titrant Volume added (mL) – Vendpoint XX. XX XX. XX XX.
XX * Trial 1 was preceded with the scout titration (trial 0). The results from the scout titration are not included in this table since they are not quantitatively accurate. The data in Table 2 were also shown to be highly reproducible. The subsequent calculations follow exactly the same steps as described above in the ‘Standardization’ section. The balance condition at the end point of the titration was then used to determine the concentration of the unknown, Cunknown, according to the following formula, where CNaOH was as shown in the ‘Standardization’ Section, Vunknown = XX.
XX mL, and Vendpoint was as shown in Table 2 above: CALCULATION/FORMULA SHOWN HERE, COMPLETE WITH UNITS The final result of this calculation, which is indeed the overall goal of this study, was equal to Cunknown = X. XX ± X. XX M @ 95% CL. Control Questions To more completely discuss the merits of this exercise and the underlying analytical techniques, the laboratory manual posed specific questions, which will now be addressed explicitly: (a) Why it is necessary to measure volumes very carefully when preparing the KHP solution (solution ), but NOT necessary to measure volumes too carefully when preparing the dilute NaOH solution (solution 4)? A complete answer to this question should be written here. The best answers will refer to specific procedural details to compare and contrast how each solution was prepared and used throughout the procedure. If deemed necessary, a student may refer to his/her actual data (i. e. volumes, weights, etc) in the response, as long as it is done unambiguously and any notations are kept constant throughout. b) Why in step 2 should the pipet and beaker be rinsed with the KHP solution (not with distilled water), whereas the Erlenmeyer flasks should be rinsed with distilled water (not with the KHP solution)? Avoid a simple, cursory explanation here. A complete answer will again refer to specific procedural details. Think deeply about the basics of a titration experiment and what values must be measured with the greatest accuracy. What if the procedure was reversed – the pipet and beaker were rinsed with DI water, while the Erlenmeyer flask was rinsed with KHP? What outcomes of the experiment would be most grossly affected or compromised?
If an anomaly like this did happen, would the goal of your experiment have been achieved successfully? (c) Give at least two reasons why it is better to read the initial volume of a buret than to adjust the volume to some round value such as zero. Your answer should be written here. Clearly explain what you mean. Don’t allow your TA to wonder what you might have meant! NOTE: THESE QUESTIONS NEED NOT BE ANSWERED IN THIS FORMAT. THEY MAY ALSO BE DONE IN A MORE NARRATIVE PARAGRAPH FORM. HOWEVER, IF THIS IS DONE, MAKE SURE IT IS CLEAR AS TO WHICH QUESTION YOU ARE ANSWERING.
NEATNESS AND ORGANIZATION ALWAYS COUNT (AND THE PERSON GRADING YOUR REPORT MAY DEDUCT ADDITIONAL POINTS IF HE/SHE HAS TO SEARCH FOR YOUR ANSWERS)!! Conclusions Re-state the final outcome of your experiment, including actual data and numbers. Sum up what you learned about the experiment and assess the reliabilility of the method. If you think there might have been better way(s) to achieve the final goals of the experiment, you may mention them here. All conclusion sections should also include possible sources of error, and how these errors may have affected your final results (concentration too high, too low, etc. ALWAYS KEEP IN MIND THE DIFFERENCE BETWEEN A MISTAKE AND AN ERROR! “A mistake is a measurement which is known to be incorrect due to carelessness, accidents, or the ineptitude of the experimenter. It’s important to distinguish mistakes from errors: mistakes can be avoided. Errors can be minimized but not entirely avoided, because they are part of the process of measurement. Data that is mistaken should be discarded. Data that contains errors can be useful, if the sizes of the errors can be estimated. ”

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