I. INTRODUCTION One of the most often used methods for analyzing acid-base equilibria is electrometric pH determination. With a glass electrode and a reference potential or a combination electrode, the pH of the sample is measured electrometrically. A series of reference solutions with known pH values are used to calibrate the measurement instrument. Another way of determining the pH is by using the dissociation constant of an acid (Ka). Note that if Ka is much greater than 1, the acid is mostly dissociated and so is said to be a strong acid. If Ka is much less than 1, the acid is dissociated only to a small extent and so is said to be a weak acid. A way to get the Ka is to get the equilibrium concentrations of HA (weak acid) and the A- (conjugate base) in the reaction. The amount of HA in the solution is determined by subtracting the amount of HA lost during ionization from the stoichiometric concentration of HA (the starting HA) in the solution. Although the amount of HA lost by ionization is substantially lower for acetic acid and most other weak acids than the initial concentration of HA in solution, the equilibrium concentration of HA is only marginally below the stoichiometric concentration. The objectives of this experiment are (1) to determine the dissociation constant of an acid in buffered and unbuffered solutions (2) to determine whether the measurements made in both solutions are equally reliable. I. METHODOLOGY A. Chemicals and Apparatus The chemicals used in the experiment are as follows: 0.20 M acetic acid, 0.20 M sodium acetate, and pH 4 & 7 buffer solutions for standardizing pH meter. The apparatus used were: two 10-mL measuring pipets with aspirator bulb (different bulbs for each pipet), six 50-mL volumetric flasks, squirt bottle with distilled water for diluting, 250-mL beaker for the rinsing of electrode, three 50-mL beaker, and a pH meter (Cyberscan PC 300). B. Procedure The pH meter was standardized with buffer solutions supplied. The meter was set to pH mode before standardizing. A buffer of pH 7 was prepared, the electrode of the pH meter was carefully submerged in the buffer, allowing it to stabilize and hold. The electrode was removed from the buffer and was rinsed with distilled water. The electrode was wiped carefully with tissue to remove excess water. The process was repeated for the buffer of pH 4. Now, the pH meter has been standardized. Determination of Dissociation Constant of an Acid in Unbuffered Solutions 20-mL of the acetic acid was prepared using the 10-mL measuring pipet and was transferred to three different 50-mL beakers (Due to the comfortability of the user, a 10mL measuring pipet was utilized instead of a 20-mL one). The pH of each beaker was measured using the pH meter following the same process used in standardizing the meter. The average pH was then calculated and the [H+] of the acetic acid was determined, this was then used to calculate the dissociation constant (Ka). The following acid sample were prepared using a 10-mL measuring pipet, squirt bottle with distilled water, and three 50-mL volumetric flask: Sample A: 20-mL acetic acid, dilute to mark with distilled water Sample B: 10-mL acetic acid, dilute to mark with distilled water Sample C: 5-mL acetic acid, dilute to mark with distilled water Note that the acid used in the samples were the acid in which the pH was measured. The molarity of each sample was determined. Each of the sample were transferred to three different 50-mL beakers and the pH was measured again using the pH meter. The Ka was then calculated for each of the three sample. Determination of Dissociation Constant of an Acid in Buffered Solutions The following samples were prepared using two 10-mL measuring pipet, squirt bottle with distilled water, and three 50-mL volumetric flask: Sample D: 20-mL acetic acid + 4-mL 0.20 M sodium acetate, dilute to mark with distilled water Sample E: 20-mL acetic acid + 20-mL 0.20 M sodium acetate, dilute to mark with distilled water Sample F: 4-mL acetic acid + 20-mL 0.20 M sodium acetate, dilute to mark with distilled water The new concentrations of the acetic acid and sodium acetate was then calculated for each sample. The samples are then transferred to three different 50-mL beakers and the pH was measured using a pH meter. The Ka was then calculated for each of the three sample. C. Mathematical equations The [H+] was calculated using the equation [H+] = 10−pH. The molarity was calculated using the equation C1V1=C2V2. The formula π²π = [π―+ ][π¨− ] [π―π¨] was used in calculating the dissociation constant of the weak acid in unbuffered solutions. The Henderson-Hasselbalch [π¨− ] equation ππ― = ππ²π + π₯π¨π [π―π¨] was used to relate the pH or the dissociation constant in buffered solutions. To derive the Ka from the Henderson-Hasselbalch equation, the pKa was calculated, then the equation Ka= 10-pKa was utilized to determine the Ka. II. RESULTS AND DISCUSSION The data in determining the average pH, [H+], and Ka of the acetic acid is summarized in Table 1 below. Table 1. Dissociation constant (Ka) of 0.20 M acetic acid Beaker Volume of acetic (mL) pH Average [H+] Ka 0.0035 6.12 x 10-5 pH 1 20 2.51 2 20 2.46 3 20 2.40 2.46 The pH recorded in the three beakers are obtained from the pH meter and the average was calculated. The literature value for Ka of acetic acid is 1.76 x 10-5. Based from this value, a 0.20 M acetic acid should have a [H+] concentration of 0.00188 and pH of 2.73. This is evident just by looking at the calculated Ka, there is a huge difference between the literature value and the calculated value. This error may be due to the purity of the stock solution or an instrumental error. Table 1 calculation: π©πππππ π ππ― + π©πππππ π ππ― + π©πππππ π ππ― = πππππππ ππ― π π. ππ + π. ππ + π. ππ = πππππππ ππ― π π. ππ = πππππππ ππ― ππ― = − π₯π¨π [π―+ ] [π―+ ] = ππ−ππ― [π―+ ] = ππ−π.ππ [π―+ ] = π. ππππππππππππ ≈ π. ππππ HA + H2O β H+ + AI 0.2 0 0 C 0.2-x +x +x E 0.19652 0.0035 0.0035 [π―+ ][π¨− ] π²π = [π―π¨] π²π = [π. ππππ][π. ππππ] [π. πππππ] π²π = π. ππ π ππ−π * Note that during calculations using ICE table, the values must not be rounded off, only the final answers should be rounded off. The calculated and measured results for the determination of Ka in unbuffered solution is shown in the table below. Table 2. Dissociation constant (Ka) of an acid in unbuffered solution Acid Volume of acetic acid pH Molarity [H+] Ka Sample (mL) (M) A 20 2.63 0.08 0.0023 7.08 x 10-5 B 10 2.79 0.04 0.0016 6.85 x 10-5 C 5 2.99 0.02 0.0010 5.52 x 10-5 Each acid sample had varying amount of acetic acid but they are all diluted to 50-mL of distilled water. Because of this they also have varying pH, molarity, [H+], and Ka. The calculated Ka for the three samples is still far from the literature value, with acid sample C the closer one. For a molarity of 0.02 M in acid sample C, the [H+] should be 0.000593 with a pH of 3.23, this is attained using the literature value of Ka for acetic acid. The error still could be from the purity of the stock solution and instrumental error. Sample calculation for Table 2 (Sample A): C1V1 = C2V2 π. ππ π΄ (ππππ³) = πͺπ (ππππ³) π. ππ π΄ (ππππ³) = πͺπ ππππ³ π. ππ π΄ = πͺπ [π―+ ] = ππ−ππ― [π―+ ] = ππ−π.ππ [π―+ ] = π. ππππππππππππ ≈ π. ππππ HA + H2O β H+ + AI 0.8 0 0 C 0.8-x +x +x E 0.0777 π²π = π²π = 0.0023 0.0023 [π―+ ][π¨− ] [π―π¨] [π. ππππ][π. ππππ] [π. ππππ] π²π = π. ππ π ππ−π The calculated and measured results for the determination of Ka in buffered solution is presented in the table below. Table 3. Dissociation constant (Ka) of an acid in buffered solution Acid Volume of Volume of Molarity of Molarity Sample acetic acid sodium the acetic of sodium (mL) acetate acid (M) acetate (mL) pH Ka (M) D 20 4 0.08 0.016 3.77 3.44 x 10-5 E 20 20 0.08 0.08 4.49 3.24 x 10-5 F 4 20 0.016 0.08 5.24 2.88 x 10-5 This time a buffer is added, the salt of the acid which is sodium acetate (CH3COONa). Varying volume is still observed in this section, both for the sodium acetate and acetic acid but they are all still diluted to 50-mL with distilled water. This will also affect the molarity, pH, [H+], and Ka of the acid sample. To compare the calculated Ka to the literature value (Ka = 1.76 x 10-5), acid sample F is the closer one. Sample calculation for Table 3 (Sample D): Molarity of acetic acid Molarity of Sodium acetate C1V1 = C2V2 C1V1 = C2V2 π. π π΄ (ππππ³) = πͺπ (ππππ³) π. π π΄ (πππ³) = πͺπ (ππππ³) π. ππ π΄ (ππππ³) = πͺπ ππππ³ π. ππ π΄ (πππ³) = πͺπ ππππ³ π. ππ π΄ = πͺπ π. πππ π΄ = πͺπ [π―+ ] = ππ−ππ― [π―+ ] = ππ−π.ππ [π―+ ] = π. πππππππππππππ ≈ π. πππππ CH3COOH + H2O β CH3COO+ I 0.08 0.016 C 0.08-x 0.016+x E 0.07983 0.016169824 [πͺπ―π πͺπΆπΆ− ] ππ― = ππ²π + π₯π¨π [πͺπ―π πͺπΆπΆπ―] π. ππ = ππ²π + π₯π¨π [π. πππππππππ] [π. πππππ] π. ππ = ππ²π − π. πππππππππ π. πππππππππ + π. ππ = ππ²π π. πππππππππ = ππ²π π²π = ππ−ππ²π π²π = ππ−π.πππππππππ π²π = π. ππ π ππ−π H+ III. CONCLUSION AND RECOMMENDATION To sum it all up, the Ka in unbuffered solution compared to the buffered solution is much higher, and compare to the literature value for Ka of acetic acid, buffered solutions are much closer than the unbuffered ones. This result tells us that measurements made in buffered solutions are much more reliable than unbuffered solutions. This is true due to the common ion effect. The common ion effect happens when an ion is added to a mixture that already contains the provided ion and is at equilibrium. When this occurs, the equilibrium changes so that less of that ion is formed. As CH3COONa dissociates it forms CH3COO- and Na+. For the partial dissociation of acetic acid, CH3COO- is also formed and this is the common ion, because of this the reaction will shift to the left therefore decreasing the [H+] and increasing the pH. It is important to calibrate instruments from time to time, as this will greatly affect the data obtained. It is also important to ensure the purity of the substance that will be used prior to the experiment. IV. REFERENCES Kotz, J.C., Treichel Jr, P.M., Townsend, J.R. (2012). Chemistry and Chemical Reactivity. Belmont, CA: Brooks/Cole, Cengage Learning Silberberg, M.S. (2009). Chemistry: The Molecular Nature of Matter and Change. Boston, MA: McGraw Hill. Zumdahl S. S. & Zumdahl S. A. (2007). Chemistry (5th ed. instructor's annotated). 222 Berkeley Street, Boston, MA: Houghton Mifflin. E9 chemical equilibria: Ka - chemvision.net. ChemVision. (n.d.). Retrieved February 22, 2023, from http://chemvision.net/102_Lab_7_KaWeakAcid_Sp18.pdf
