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Chemistry Lab 2 Determintion of Ka Titration

by: Pooja Dave

Chemistry Lab 2 Determintion of Ka Titration CHM 113

Marketplace > University of Miami > Chemistry > CHM 113 > Chemistry Lab 2 Determintion of Ka Titration
Pooja Dave
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This is the lab report for the second full length report required by CHM 113
General Chemistry 1 lab
Dr. Tegan Eve
Class Notes
General Chemistry
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This 3 page Class Notes was uploaded by Pooja Dave on Thursday September 29, 2016. The Class Notes belongs to CHM 113 at University of Miami taught by Dr. Tegan Eve in Fall 2015. Since its upload, it has received 4 views. For similar materials see General Chemistry 1 lab in Chemistry at University of Miami.


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Date Created: 09/29/16
Pooja Dave CHM 113, Section FY Determination of Ka: Titration of a Weak Acid Introduction: The Ka value, the acid dissociation constant, represents the strength of an acid. The lower the Ka, the weaker the acid. There are two methods of finding the Ka value, both involving the usage of pH to find pKa. The first method is to measure the pH of a solution containing a known concentration of a weak acid. From there, the pKa may be calculated using the Henderson-Hasselbach equation. The second method is to measure the pH at half-neutralization point in titration of a weak acid with a strong base. Because the pH at half-neutralization point is equal to the pKa, the Ka can then be easily calculated. Titration of the acid by the strong base allows the unknown concentration of the weak acid and the pH to be measured for both methods. The equivalence point is found by looking at the point on the graph where the slope changes from increasing to decreasing. This means that it is also the inflection point so that first and second derivatives can be used in its calculation. In this experiment, we use acetic acid (HC₂H₃O₂) to show how the Ka value and pH are related. Procedure: 1. Obtain 20mL of weak acid solution in a 100mL beaker 2. Begin calibrating the drop counter using distilled water and LoggerPro 3. Fill the reagent reservoir with about 60mL NaOH, letting a little bit out so tip is filled with NaOH 4. Place the 20mL of weak acid under the drop counter 5. Insert the pH sensor and magnetic stirring bar inside the beaker and begin data collection 6. Continue data collection until the graph shows a large increase in pH 7. Changing the y axis to show second derivative, find the highest point on the graph to find the pH at the equivalence point 8. Divide the volume added at equivalence point by half to find the pH at half- neutralization point Equations: HC₂H₃O₂ + NaOH -> HOH + NaC₂H₃O₂ (Molarity of base) * (volume at equivalence point) = (molarity of weak acid) * (volume of acid originally) [ A⁻] pH = pKa + log ( [HA] ) pH½ = pKa pKa = -log(Ka) Observations: During our experiment, the greatest error was in calibrating the drop counter. Some drops were not counted properly by the LabPro computer interface, making the calibration of the counter inaccurate for recording the data. Data: Volume at Equivalence Point from Titration – 38.44mL Molarity of NaOH – 0.100M Molarity of Weak Acid originally - 0.1922M Volume at half-neutralization point – 19.25mL pH at Equivalence Point – 9.90 pH at half-neutralization point – 5.15 Ka from pH at Equivalence Point – 6.55 ×10−11 Ka from pH at half-neutralization point – 7.08 ×10 −6 Calculations: .1M )×(38.44mL )=[HA ]×(20mL ) [HA] original = 0.1922M 0.1M 0.1922M (¿) 9.90=pKa+log¿ pKa = 10.182 10.182 = -log(Ka) ×10 −11 Ka from equivalence point = 6.55 5.15 = pKa 5.15 = -log(Ka) Ka from half-neutralization point = 7.08 ×10 −6 Discussion and Conclusions: The equivalence point obtained from titration shows when the weak acid is neutralized by sodium hydroxide base, or the inflection point of the curve. On the graph created by collecting data, the equivalence point is contained in the nearly vertical portion. The point is half-way up the linear portion of the titration curve and can be calculated using the derivatives. The first derivative peaks at the equivalence point, while the second derivative passes through zero at this point. The volume and pH at this point can be used in determining the original concentration of the weak acid along with the Ka value by the Henderson- Hasselbach equation. Our lab shows that the equivalence point is at 38.44mL with a pH of 9.90. Another method is to use the half-neutralization point in calculating the Ka value. The half-neutralization point is halfway to the equivalence point. In our lab, this point was at approximately 19.25mL with a pH of 5.15. The Ka value at the equivalence point is 6.55e-11 while at the half-neutralization point it is 7.08e-6. The difference between the two values illustrates the reliability of the half-neutralization point method, as the actual Ka value for acetic acid is 1.77e-5. Errors:  The calibration of the drop counter may have been inaccurate. This would affect the equivalence point and the concentration of weak acid and Ka, making them lower.  The calculation of the half-neutralization point may have been wrong. Dividing the equivalence point by half may have given us a larger pH and a higher Ka value than in reality.  The collection by the pH sensor may not have been accurate due to some interference with the sensor and the distilled water solution it was placed in. This would alter the data collected, making the pH level higher than it actually was.  More NaOH may have entered the beaker than was counted, making the pH higher than it actually was. This then would have made the Ka higher. Graph:


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