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Potentiometric Titration of Sodium Carbonate
Otieno O. Victor
University of Detroit Mercy
Quantitative Analysis Lab CHM 3880
Fall 2011
Partner: Edwin Gay

Abstract
The PH at each point during the titration of sodium carbonate unknown sample was determined. An Unknown sample of Na2CO3 was titrated with a standard HCL solution. In addition to titration, the pH at each point of titration was measured using PH meter. The % of the unknown Na2CO3 was 25.83% Introduction1

The purpose of this experiment was to determine the actual pH at each point during the titration of a sodium carbonate unknown with the use of pH meter. An acid-base titration is a procedure which is used to determine the concentration of an acid or base. A measured volume of an acid or base of known concentration is reacted with a sample to the equivalence point. However, there are difficulties in completing acid-base reactions with the aid of visual indicators. This is likely due to factors like unsuitable color change for a given type of titration. Also, this might be due to student’s being color-blind to certain indicator color changes. Moreover, certain solutions might be already be colored. To limit such uncertainties and difficulties, potentiometric titration is applicable. Potentiometric titration can be used to distinguish acids. The amount of titrant added helps determines the pH of a solution. Usually, a change in pH observed is small. However, at endpoint, the pH observed is often of a sharp change. The sharpness of the change is reflective of the strength of the acid or base. To determine the end point, all the point is graphed and then location of a sudden change in pH determined. The equations used were: Equation: Na2CO3 + HCl ↔ H+ + NaCO3- for first equivalence point reaction. Equation 2: Na2CO3 + HCl ↔ H+ + NaCO3 for second equivalence point reaction Experimental procedure1

In standardization of HCL, a dried and cooled Na2CO3 was placed in a weighing bottle. Then, 3 samples of primary standard Na2CO3 were weighed in 125 mL Erlenmeyer flask. The masses were 0.1833g, 0.1835g, and 0.1834g. Then, 50mL of distilled water was added to each flask. Then, 3 drops of phenolphthalein indicator was added to each flask. Then, 50mL burette was filled with standard HCl solution. The solutions in each were titrated to phenolphthalein end point. Then, 3 drops of bromo-cresol green indicator was added. Then, titration was continued until bromo-cresol green indicator begun to turn green. The solutions were then boiled for 10 minutes and titrated to bromo-cresol green end point. In titration of unknown, a sample of 0.3959g was weighed, placed in 250-mL beaker and 50mL of distilled water added to the beaker by pipette. The solution was then stirred with glass rode. The PH meter was standardized with 2 buffer solutions. The electrodes were washed and gently wiped with soft tissue. The electrodes were immersed in the solution and left immersed for the entire period of titration. The starting PH was recorded. A burette was filled with the standard HCl solution and starting volume noted. About 1.0mL of HCL was added to the solution and thoroughly stirred and PH recorded. The titrant addition was continued at about 1.0mL at a time. The PH and volume changes were recorded. At about 1.0-1.5mL of first equivalence, the titratnt was added in an increment of 0.2mL. The burette and the PH were recorded. However, me and my lab partner forgot to boil the solution when we were about 1.0 of second equivalence point. We continued with the titration with small increment of titrant until about 1.0-2.0mL beyond second equivalence point. Then, the electrodes were removed and rinsed thoroughly and then re-immersed in distilled water. Results

The table1 shows the titration results for the known mass of Na2CO3 Table1: titration results
Trial #| Mass of Na2CO3 (g)| Volume of 1nd equivalence(mL)| Volume of 2nd equivalence(mL)| Last amount to 2nd equivalence (mL)| Total...
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