Idea Transcript
Determination of the Ionization Constant of a Weak Acid Determination of Ka for a weak acid in water: A weak acid, HA, dissociates into H+ ions and the conjugate base, A- ions, as shown in the following equilibrium equation: HA Ö H+ + A- Equation (1) The equilibrium constant, sometimes called ionization constant or dissociation constant, is:
Equation (2) If we have only the acid in water, then equal amounts of H+ and A- are formed, so: [H+] = [A-] Equation (3) If we can determine the concentration of [H+] and the concentration of [HA], we can calculate Ka. In this experiment, we will titrate the acid with a standard NaOH solution to determine [HA], and we will measure the pH of the original acid solution to determine [H+]. Determination of Ka from a buffer solution: As NaOH is added in a titration, the concentrations of [H+] and [HA] will go down, and the concentration of [A-] will go up. Thus Ka is the product of [H+] multiplied by the ratio of [A-]/[HA]. Again, we will determine the concentration of [H+] by measuring the pH. We can adjust the ratio of [A-]/[HA] by varying the amount of NaOH added. In the experiment, we will determine Ka three times by varying the ratio. A buffer solution exists at any point during the titration, except for the original acid and the equivalence point and beyond. All the points between the beginning and end point have appreciable mounts of both weak acid and conjugate base, and such a solution is called a buffer solution. Determination of Ka at the equivalence point of a titration: The titration of a monoprotic weak acid to the equivalence point with a strong base (NaOH) results in the formation a conjugate base, A- and water. HA + OH- Ö H2 O + A- Equation (4) At the equivalence point, the solution will be basic since the conjugate base hydrolyzes in water to form hydroxide ion. The equilibrium equation for the hydrolysis of the conjugate base A- is the opposite of the above titration equation: A- + H2 O Ö HA + OH- Equation (5) The [OH-] and the [H+] can be determined from the pH, and [HA] = [OH-]. Nearly all of the HA was converted to A- at the equivalence point, so [A-] is calculated by taking the total amount of HA titrated (the same as the amount of NaOH added) and dividing by the total volume at the equivalence point. This gives you the information you need to calculate Ka. There is a drawback to this method. At the equivalence point, the pH is changing rapidly, so that the precise pH at the equivalence point must be obtained by averaging the pH before and after the equivalence point. (This method will not work at all for the first ionization constant of a diprotic acid.) Experimental Procedure: Each student will be given a sample of a weak monoprotic add whose constant is to be determined. First measure the pH of this acid as precisely as possible, using a pH meter if available. (If a pH meter is not available, use universal indicator paper, but this method is much less accurate.) Be careful in handling the electrode for the pH meter, as they are fragile and can be easily damaged or broken. Rinse the electrodes before and after, using deionized water. When not in use, the electrodes must be stored in water or a buffer solution. Precisely measure 20.00 mL of the unknown acid into an Erlenmeyer flask using a 20-mL pipette or a 50-mL burette. Add 2 drops of phenolphthalein indicator (unless a different indicator is recommended by the instructor). Titrate using a standard 0.1 M solution of NaOH added from a 50-mL burette (The phenolphthalein indicator turns barely pink at the end point, and remains pink for at least 20 seconds, on stirring.) Record the precise molarity of the NaOH (4 significant figures) the precise volume of unknown acid (nearest .01 mL) and the precise volume of NaOH needed to titrate to the equivalence point (nearest .01 mL). Repeat to verify the result. Divide the volume of NaOH by 4. This is the volume of NaOH that you will add for each of the following tests: Now start with a new 20.00 mL portion of the unknown acid. Add ¼ the volume of NaOH that you needed for the complete titration. Measure the pH of the solution and calculate the [H+]. One-fourth of the HA has been converted to A-. This means that ¾ of the HA remains. The ratio of [A-]/[HA] = 1 /3 = .3333 = ¼ divided by ¾. From Equation (2), Ka = .3333 [H+] To this same solution, add an additional 1/4 volume of NaOH. Measure the pH and calculate the [H+]. You are now half-way to the equivalence point, so half the HA has been converted to A- Therefore [HA] = [A-], and the ratio [A-]/[HA] = 1. Ka = [H+ ],or pH = pKa. Note that in any titration of a week add with strong base, the pH half-way to do end-point equals the pKa . (It can also be shown that in the titration of a weak base with a strong acid, the pOH half-way to the end-point equals the pKb of the base.) Next, to the same solution add an additional ¼ volume of NaOH. Measure the pH and calculate the [H+]. You are now ¾ of the way to the equivalence point, so ¾ of the HA has been converted to A- Therefore the ratio of [A-]/[HA] = ¾ divided by ¼ = 3, and Ka = 3 [H+]. If you add another ¼ volume of NaOH, you are now at the equivalence point. However, the pH measured at this point is not very accurate, because the curve is at a point where the pH is rising rapidly. If you add still another ¼ volume of NaOH, you are now somewhat past the equivalence point. Measure the pH and calculate the average of this pH and the pH at the ¾ point (before the equivalence point). This average is a better value for the pH at the equivalence point. Calculate the [H+], and the [OH-] concentration from this average pH. [A-] = MNaOHx 20/total volume of the solution. Using equation (2), calculate Ka. Graph: On linear graph paper, plot pH (y-axis) vs. mL of NaOH (x-axis). You should have a total of 6 points to plot. Turn in your graph with the report form.
Determination of the Ionization Constant of a Weak Acid Student's Name___________________________ Unknown Sample Number ____________________ Observed pH of unknown ____________________ Titration data: Volume of unknown acid ______________ ______________ Volume of __________M NaOH ______________ ______________ Calculated Normality of acid ______________ ______________ Partial neutralization mixtures: ACID BASE RATIO PH [H+] (mL) (mL) [A-]/[HA] Mixture 1 1/4 of the way) _________ _________ _________ _________ _________ Mixture 2 (half-way) _________ _________ _________ _________ _________ Mixture 3 (3/4 of the way) _________ _________ _________ _________ _________ Mixture 4 (equivalence point) _________ _________ XXXXXX _________ _________ Mixture 5 (past equivalence point) _________ _________ _________ _________ _________ CALCULATION of Ka (Clearly show method of calculation for each determination.): Mixture 1: Ka = Mixture 2: Ka = Mixture 3: Ka = Mixture 4: (Hint- first calculate [A-]; also [HA] = [OH-] (based on pH of mixture 4, less reliable) Ka = pH at equivalence point based an Mixture 3 and Mixture 5 ________ [H+] = _____________ [A-] = same as mixture 4 [HA] = [OH-] =_____________ Ka =