Chemistry

AcidpH measuredH3OH3O theoreticalpH theoreticalerrorAve error0.1 M HOAc2.811.55E-0032.881.31E-0032.434.970.01 M HCl2.157.08E-00321.00E-0027.5Table 1. Instrumental error of the pH meter

Trial
mL NaOH used
pHAssuming salt  acidAssuming salt  acidpKaKaKa avepKaKaKa aveFor 0.1 M acetic acid127.34.664.622.39E-0052.41E-0054.662.19E-0052.24E-005226.54.644.612.43E-0054.642.29E-005For 0.1 M chloroacetic acid128.12.752.72.00E-0032.58E-0032.751.78E-0032.30E-003228.12.552.53.17E-0032.552.82E-003For 0.1 M dichloroacetic acid132.71.81.682.07E-0022.25E-0021.81.58E-0021.72E-002232.71.731.612.44E-0021.731.86E-002For 0.1 M trichloroacetic acid125.81.531.523.05E-0023.01E-0021.532.95E-0022.92E-002225.81.541.532.98E-0021.542.88E-002Table 2. Values for pKa and Ka of acetic acid and several other acids

Anilinium hydrochloride concentration, MpHKh0.152.182.94E-0044.40E-0020.082.75.36E-0052.66E-002Table 3. Values for Kh and  for anilinium chloride at different concentrations

DISCUSSION
In Table 2 are presented two values for the acid dissociation constant Ka of acetic acid and other acids at 22.5 C, using two assumptions. Assuming that the concentration of the salt is not equal to the concentration of the added acid utilizes the data on the amount of standard NaOH solution used to neutralize the acid, seeing as the volume of used NaOH is not equal to 25 mL. Assuming that the concentration of salt and added acid are equal disregards the volume of NaOH used. Of the two assumptions, the first assumption seems more likely the second assumption is included here for comparison. For both assumptions, the Henderson-Hasselbach equation pH  pKa  log saltacid was used to arrive at the value of pKa, which is the negative logarithm of the Ka value. From the literature (Atkins et. al 2004, p. A-8 to A-9), the value of Ka at 25 C for acetic acid, chloroacetic acid, dichloroacetic acid and trichloroacetic acid are 1.75E-005, 1.36E-003, 3.31E-002 and 3, respectively. Comparing these values with the measured values in Table 2, it can be computed that the percent errors of the determination are, ignoring the difference in temperature 37.64, 89.95, 31.92 and 99, respectively, assuming salt  acid and 27.96, 68.99, 47.95 and 99.03, respectively, assuming salt  acid.

It will be noticed that the measured value of 3.01E-002 as Ka for trichloroacetic acid is far smaller than the literature value of 3. The high literature value of Ka indicates that trichloroacetic acid behaves as a strong acid, and so dissociation is virtually complete. The introduction portion of the manual implies that the procedure is based on weak-acid equilibria. But then equilibrium is essentially not established with trichloroacetic acid, as it is completely dissociated, so the method for determining Ka used here is not applicable to this acid.

The literature values indicate that Ka and thus acidity increases as the acid becomes more and more chlorinated. This can be explained in terms of the Lewis acidbase theory the presence of chloride draws electrons away from the acidic hydrogen side, making the hydrogen electron-deficient and thus more acidic. The more chlorinated the acid molecule is, the more-electron deficient and the more acidic the hydrogen is.
Buffer solutions are solutions containing an acidic or basic species and its conjugate in equilibrium with each other. Buffer solutions serve to maintain the pH of a solution at a constant value when small amounts of acid or base are added to it, or when it is diluted. (Skoog et.al 2004, p. 251)

The measured values of the hydrolysis constant Kh and fraction  of anilinium chloride dissociated are computed from the measured pH by taking note that   (Ka  C) and, from the expression for Ka, anilineanilinium    Ka  H3O. Combining these two expressions give the value of Ka for each concentration, and from the value of Ka, Kh and  can be computed. The measured value of Kh for 0.15 M anilinium chloride and 0.075 M anilinium chloride are 2.94E-004 and 5.36E-005, respectively, while for  the values are 4.40E-002 and 2.66E-002, respectively. From the literature value of Kw at 298 K, Kh is computed to be 2.51E-005. The percent difference of the theoretical values to the measured values are then computed to be 1071 and 113 respectively.

The values for Kh and  are expected to decrease with increasing concentration. At infinite dilution, weak electrolytes are said to be completely dissociated, but as the concentration increases, the presence of other electrolyte molecules drive the equilibrium in favor of the undissociated species. However, the opposite is observed in the results of this experiment. This student cannot find an explanation why this is so.

Using the literature value of Ka for acetic acid at 25 C and assuming that hydrochloric acid dissociates completely, the theoretical pH of 0.1 M acetic acid and 0.01 M hydrochloric acid are computed to be 2.88 and 2.00 respectively. The percent difference between the theoretical value and the measured value is then computed to be 2.43 and 7.5, respectively.