Results and DiscussionMany titrations are now done automatically by instruments which are calibrated to deliver small increments of titrant between pH measurements. Titration is one of the techniques by which the characteristics of any substance can be analyzed. Potentiometric titration is one of them. During titration, pH were determined and plotted against the buret readings to obtain characteristic curves that were used to determine the equivalence point/s of the solution [3].In this experiment a weak acid as sample (coke) and a strong base (NaOH), sodium hydroxide is commonly known as lye or caustic soda and dissolving NaOH in water generates considerable heat, was used where three trials made. See tratment of data (tables) in Appendix A. Data from trial 3 were used thoughout the discussion. The amount of NaOH used to prepare 1M NaOH was 0.4121 grams (see calculation in Appendix B). By reading the pH meter, the pH acquired by the sample is 4.01. In second year college, students were familiarized with the appearance of a plot of pH vs. mL of base added in typical titraion. Using these knowledge, the behavior of the pH at the beginning, at the neuralization and towards the end of titration were compared with the titration graphs attained from the experiment using potentiometric titration. Figure 1a and 1b (see reference 4) were used to compare the graph obtained from the experiment.

Figure 1a. Titration curve for weak acid (ethanoic acid) vs. strong base (sodium hydroxide). Running acid into alkali. For the first part of the graph (Figure 1a), there is an excess of sodium hydroxide. The curve will be exactly the same when an addition of hydrochloric acid to sodium hydroxide was made. Once the acid is in excess, there will be a difference. Past the equivalence point, a buffer solution containing sodium ethanoate and ethanoic acid was gotten. This resists any large fall in pH.

Figure 1b. Titration curve for weak acid (ethanoic acid) vs. strong base (sodium hydroxide). Running alkali into acid

The start of the graph shows a relatively rapid rise in pH but this slows down as a buffer solution containing ethanoic acid and sodium ethanoate is produced. Beyond the equivalence point (when the sodium hydroxide is in excess) the curve is just the same as that end of the titration graph of strong acid vs. strong base.

Equi pt

Figure 2. Titration curve from experimental potentiometric titration. (pH data versus the volume readings of the base. On the graph shown above the weak acid only partially dissociates from its salt. The pH raised normally at first, but as it reaches a zone where the solution seems to be buffered, the slope levels out. After this zone, the pH rises sharply through its equivalence point and levels out again like the strong acid/strong base reaction. There are two main points to notice about this curve. The first is the half-equivalence point. This point occurs halfway through a buffered region where the pH barely changes for a lot of base added. The half-equivalence point is w hen just enough base is added for half of the acid to be converted to the conjugate base. When this happens, the concentration of H+ ions equals the Ka value of the acid. Take this one step further, pH = pKa. The second point is the higher equivalence point. Once the acid has been neutralized, notice the point is above pH=7. When a weak acid is neutralized, the solution that remains is basic because of the acid's conjugate base remains in solution.V/mLpHVmidph/vol

04.01

14.460.50.45

1.24.91.12.2

1.35.41.255

1.456.11.3754.666666667

1.66.771.5254.466666667

1.77.351.655.8

1.87.831.754.8

2.28.3321.25

2.68.842.41.275

39.382.81.35

3.29.923.12.7

3.510.453.351.766666667

3.610.963.555.1

4.111.463.851

4.311.744.21.4

This is perhaps the simplest method for interpreting pH titration data. Because experiments do not always proceed so neatly, various other methods are often used to determine equivalence points. For example, since the region of maximum slope contains this point, a graph of the slope (known as a "derivative" in calculus terminology) versus the average volume for the slope interval provides a way to zero in on the desired volume. The slope could be represented as pH/V (i.e., y/x) and Vmid. A plot of the same data as shown earlier was treated in this way is shown below:

Figure 3. First derivative. The equivalence point corresponds to the volume at the peak of the curve

Occasionally it is still difficult to judge the equivalence point even by this method. Sometimes the "peak" is rounded or has a plateau. A second derivative plot (the change in the slope) may be used in such cases. The values for the plot are obtained by operating on the first derivative data in the same way as the orig ginal data was processed. A plot of ratio/V2 vs. V2mid for the same titration data is shown below:

ratioV2V2midratio/V2

3.460.550.7756.290909091

0.20.11.12

-1.80.151.225-12

50.151.37533.33333333

4.60.0751.487561.33333333

-7.50.0751.5625-100

-2.566670.1251.6625-20.53333333

-1.10.3751.9125-2.933333333

3.8666670.352.27511.04761905

-2.20.152.525-14.66666667

-10.252.725-4

2.80.22.9514

-1.850.153.125-12.33333333

-0.916670.253.325-3.666666667

-0.933330.553.725-1.696969697

Figure 4. Second Derivative, The euivalence point corresponds to the volume where the curve crosses the X-axis.

The equivalence point is found at the average volume where the function crosses y = 0. The downside to the derivative methods is that each involves a compromise in the accuracy of the volume since the interval chosen for the derivative requires an average volume. Making the intervals small improves the accuracy and is a good reason for adding titrant in very small increments in the vicinity of the equivalence point. For the reaction of a weak acid, such as the acid used in this experiment, with a strong base (NaOH), the equivalence point pH is no longer 7 unlike with the reaction between strong acid and strong base. The equivalence point region on the titration curve is not centered at pH 7. There is also a more pronounced "level" region at the beginning of the titration and approaching the equivalence point as buffering occurs due to the partial neutralization of the acid as base is added, and the simultaneous formation of substantial amounts of conjugate base.

ConclusionPotentiometric titration involves the measurement of the potential of a suitable indicator electrode with respect to a reference electrode as a function of titrant volume. [5]Potentiometric titrations provide more reliable data than data from titrations that use chemical indicators and are particularly useful with colored or turbid solutions and for detecting the presence of unsuspected species. Weak acid strong base curve starts at a pH higher than 1 which is generally the case with strong acids. This is because of the incomplete dissociation of the acid. There is a partial dissociation of the acid in to H+ and A- ions. A- combines with the B+ of the base while H+ and OH- get neutralized. The AB salt formed acts as a conjugate acid and base which will assist further dissociation of weak acid. Once the equivalence point is reached the strong base affects the increase of pH rapidly.

Recommendation In this experiment our group recommend that students must have a well-trained hand because it is an important tool. Students should strive to develop good control of the buret stopcock, delivering single drops with 100% reliability and no false squirts. Accurately reading the volume on the buret is another important skill. Be sure the meniscus is at eye-level when recording a volume. Many people find it helpful to place a card behind the buret with a white/black boundary to help determine the exact position of the meniscus.

References[1] Neumann, Erzsbet. (2010). Advanced Potentiometry: Potentiometric Titrations and Their Systematic Errors.[2]Skoog, D. ,West D. ,Holler, J. ,Crouch, S. (2004). Fundamentals of Analytic Chemistry (Philippine Edition).[3]Agbayani, Virgilio. (2014). Laboratory Manual in Physical Chemistry for Rngineers 1[4] http://chemistry.about.com/od/acidsbase1/ss/titrationcurves_2.htm[5]http://memo.cgu.edu.tw/hsiu-po/Analytical%20Chem/Lecture%207.pdf

Appendix A

Table 1. Treatment of data for first trial in potentiometric titrationV/mLpHVpHVmidph/volratioV2V2midratio/V2

04.01

14.5710.560.50.56

1.14.980.10.411.054.13.540.550.7756.436363636

1.355.750.250.771.2253.08-1.020.1751.1375-5.828571429

1.46.480.050.731.37514.611.520.151.376.8

1.57.020.10.541.455.4-9.20.0751.4125-122.6666667

1.67.50.10.481.554.8-0.60.11.5-6

2.27.970.60.471.90.783333333-4.016670.351.725-11.47619048

2.48.440.20.472.32.351.5666670.42.13.916666667

2.78.970.30.532.551.766666667-0.583330.252.425-2.333333333

39.460.30.492.851.633333333-0.133330.32.7-0.444444444

3.310.020.30.563.151.8666666670.2333330.330.777777778

3.510.540.20.523.42.60.7333330.253.2752.933333333

3.711.010.20.473.62.35-0.250.23.5-1.25

4.311.510.60.540.833333333-1.516670.43.8-3.791666667

4.611.690.30.184.450.6-0.233330.454.225-0.518518519

4.6511.750.050.064.6251.20.60.1754.53753.428571429

Graphs Obtained From Table 1

Table 2. Treatment of data for second trial in potentiometric titration

V/mLpHVpHVmidph/volratioV2V2midratio/V2

04.01

14.5510.540.50.54

1.14.950.10.41.0543.460.550.7756.290909091

1.25.370.10.421.154.20.20.11.12

1.45.850.20.481.32.4-1.80.151.225-12

1.56.590.10.741.457.450.151.37533.33333333

1.557.190.050.61.525124.60.0751.487561.33333333

1.657.640.10.451.64.5-7.50.0751.5625-100

1.87.930.150.291.7251.933333-2.566670.1251.6625-20.5333333

2.48.430.60.52.10.833333-1.10.3751.9125-2.93333333

2.58.90.10.472.454.73.8666670.352.27511.04761905

2.79.40.20.52.62.5-2.20.152.525-14.6666666

39.850.30.452.851.5-10.252.725-4

3.110.280.10.433.054.32.80.22.9514

3.310.770.20.493.22.45-1.850.153.125-12.3333333

3.611.230.30.463.451.533333-0.916670.253.325-3.66666666

4.411.710.80.4840.6-0.933330.553.725-1.69696969

Graphs Obtained From Table 2

Table 3. Treatment of data for third trial in potentiometric titration

V/mLpHVpHVmidph/volratioV2V2midratio/V2

04.01

14.4610.450.50.45

1.24.90.20.441.12.21.750.60.82.916666667

1.35.40.10.51.2552.80.151.17518.66666667

1.456.10.150.71.3754.666666667-0.33330.1251.3125-2.66666666

1.66.770.150.671.5254.46666666-0.20.151.45-1.33333333

1.77.350.10.581.655.81.3333330.1251.587510.66666667

1.87.830.10.481.754.8-10.11.7-10

2.28.330.40.521.25-3.550.251.875-14.2

2.68.840.40.512.41.2750.0250.42.20.0625

39.380.40.542.81.350.0750.42.60.1875

3.29.920.20.543.12.71.350.32.954.5

3.510.450.30.533.351.76666666-0.93330.253.2253.733333333

3.610.960.10.513.555.13.33330.23.4516.66666667

4.111.460.50.53.851-4.10.33.7-13.6666666

4.311.740.20.284.21.40.40.354.0251.142857143

Graphs Obtained From Table 3

Appendix BCalculation for the preparation of 0.1M NaOH

0.399969 grams