The Titration of Acids and Bases (Lab Report Sample)

The Titration of Acids and Bases Name Lab partner’s name Course and section number Instructor’s name Date of experiment Purpose The goal of this experiment was to get familiarized with techniques of titration, the volumetric procedure of analysis, and be able to ascertain the acidic content in an unknown substance. Introduction Neutralization is defined as the reaction between a base and an acid. The reaction is used to accurately determine the amount of Hydroxide prepared. The process by which the concentration of a solution can be determined is known as standardization. To attain the acidic content in an unknown sample, the standard base required to neutralize this acid has first to be measured by use of a burette, a process known as titration (Skiba et al. 2016). The point at which an acid has neutralized a base or a base has neutralized an acid can be determined by the use of an indicator solution, which undergoes a color change once equal amounts of acid and base are present. Specifically, this color change is defined as endpoint of the titration (Aslan et al. 2014). An indicator would tend to have a different color at different pH values. For instance, in acidic solutions, phenolphthalein is colorless and pink in alkaline solutions. In this experiment, Sodium Hydroxide solution was standardized by being titrated through a sample of pure and known quantity of Potassium Hydrogen Phthalate, KHC8H4O4 (abbreviated KHP). KHP with a molar mass of 204.2 and is a monopromatic. For KHP to be neutralized, the balanced equation is as follows. 23717259271000KHC8H4O4(aq) + NaOH(aq) H2O(l) + KNaC8H4O4(aq)(1) The equivalence point for titration is when stoichiometrically equivalent amounts of two different reactants are put together. It is theoretical and can be estimated by making certain observations like color changes at an end point. At the equivalence point, moles of NaOH= moles of KHP(2) Molarity (M) can be defined as the number of moles of a solute for every liter of solution. That is, M= moles of solutevolume of solution (in liters)= moleliter also=mmolml(3) Hence M×liters=moles of solute(4) It is, therefore, possible to calculate the molarity of the measured volume of NaOH solution that can neutralize a given amount of KHP. This can then be used to determine the amount an acid, such as KHP that is in a specific weight of an impure sample. The percentage of the KHP in this sample is then as follows: % of KHP = grams of KHPmass of sample×100 Phenolphthalein is the indicator chosen because its color change is close to the equivalence point. Procedure To prepare 0.1M NaOH, 500ml of distilled water was boiled to remove CO2, which was succeeded by cooling and transferring it to a 1-pint bottle. 3 ml of carbonate solution of a stock solution of NaOH was added and vigorously shaken for more than a minute. Then, a 50 ml burette was prepared by cleaning it with soap solution then thoroughly rinsing it with tap water then by five 10ml portions of distilled water. Standardization of NaOH Solution Firstly, 400-450ml of CO2 free water was prepared by boiling it. Three samples of accurately weighed KHP (0.4-0.6g) were put into 250 ml Erlenmeyer flasks. The weights were recorded and flasks labeled accordingly. The previously prepared distilled water (100ml to each sample) was added. The flasks were gently warmed and swirled until the salt dissolved. After that, two drops of the phenolphthalein were added to each flask. The burette was rinsed with about four 5ml portions of the NaOH solution (0.1M). It was filled with the solution while ensuring that no air was in the tip by running some of the solution. Then, it was left to stand for approximately 30 seconds before ensuring the lower part of the meniscus was at the zero mark. The NaOH solution was slowly added to one of the flasks while gently swirling it. After the pink coloration that formed with addition of the solution disappeared, the NaOH was added drop wise. The flask was constantly swirled and when the solution turned pink, it was noted that the end point was reached. The procedure was repeated for the other two flasks and the burette readings were recorded for each trial. The molarity of NaOH was determined for the three tests. Analysis of an Unknown Acid In particular, this part was done for 20ml of standardized NaOH, assuming the unknown sample to be 75% KHP. Three portions of the sample were weighed and placed in three, 250 ml flasks. The sample was dissolved in 100ml of CO2 free water and 2 drops of the indicator were added. The solution was titrated to the faintest shade of pink just like in the procedure above. The percentage KHP was calculated afterwards. Data and Results Table 1 The standardization of NaOH Trial 1 Trial 2 Trial 3 Mass of KHP(g) 0.45 0.46 0.463 Final buret reading (ml) 22.6 22.56 22.5 Initial buret reading (ml) 0 0 0 NaOH used (ml) 22.6 22.56 22.5 Molarity of NaOH (mols/liter) 0.973 0.975 0.977 Molarity Trial 1= no. of moles of soluteliters of solution= mass of KHPmolar mass of KHPliters of solution= 0.45204.222.6/1000 = 0.02222.6/1000=0.973 moles Trial 2=no. of moles of soluteliters of solution= mass of KHPmolar mass of KHPliters of solution= 0.46204.222.56/1000 = 0.02222.56/1000=0.975 moles Trial 3 =no. of moles of soluteliters of solution= mass of KHPmolar mass of KHPliters of solution= 0.463204.222.5/1000 = 0.02222.5/10000.977 moles Average molarity = 0.973+0.975+0.9773=0.975mol/liter Standard deviation, =1NN=1N(x1-)2 where is summation, x1 is trials, is mean (average), and N is number of trials. Hence, =(0.973-0.975)2+(0.975-0.975)2+(0.977-0.975)23= ±0.0016 Table 2 The analysis of the amount of the unknown acid Trial 1 Trial 2 Trial 3 Mass of unknown(g) 0.418 0.428 0.438 Final buret reading (ml) 16.5 20.2 18.4 Initial buret reading 0 0 0 NaOH used (ml) 16.5 20.2 18.4 Mass of KHP in ...