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 solution 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.