Determination of Ka of several weak acids OBJECTIVE To determine the acid ionization constant, KC, for acetic acid and an unknown monoprotic acid by using indicators and by using a pH meter. CONCEPT TO BE TESTED The value of Ka, for a weak acid can be determined from the concentrations of the species present at equilibrium. TEXT REFERENCES (1) Whitten, K, W., Davis, R. E., and Peck, L. General Chemistry, 5th ed., Saunders College Publishing, Philadelphia, 1996, Sections 11.1.–11.3, Chapters 18 and 19 (2) Safety, pp1–3, (3) Laboratory Techniques Section K and L and Figure L.1. INTRODUCTION The equation, H2O(l) + HA(aq) H3O+(aq) + A–(aq), describes the equilibrium conditions when a weak acid, HA, dissolves in water and ionizes to form H3O+ and K ions. The ratio of the algebraic product of the concentrations of the ions to the concentration of the unionized acid is constant and is given the symbol, Kc. H O  A 

KA 
3
HA 
-
(Eq. 1) Kc applies to equilibrium conditions of the ionization reaction. If the acid is relatively strong, the extent of ionization is reflected in a relatively large value for Kc. A large value of Kc is the result of large concentrations of H3O+ and A– ions divided by the smaller concentration of unionized HA molecules, A weaker acid has smaller concentrations of H3O+ and A ions and a larger concentration of HA molecules. Its Kc value will be smaller. The values of Kc for different acids tell us about their relative strengths as acids. There are many important weak acids. For example, vinegar is a dilute solution of acetic acid, CH3COOH. Carbonic acid, H2CO3, is the weak acid formed by CO2 and H2O in carbonated bever‐
ages, in blood, and in many other systems. The value of Kc, for an acid can be calculated from the measured values of the concentration of the ionized and un‐ionized species present in solution at equilibrium. Example 1: The equilibrium concentrations for a given acid were found to be as follows. H3O+ HA A– Equil. conc. 1.0 x 10–1 M 1.3 x 10–3 M 1.3 x 10–3 M + H O A   1.3 x 10  1.3 x 10 

HA
1.0 x 10  
KC
-
-3
-3
3
-1
Ka = 1.7 x 10–5 The [H3O+] in an aqueous solution can be expressed as pH, which is defined as pH = ‐log [H3O+] (Eq. 2) Equilibrium constants can also be expressed in a similar way. pKa = ‐ [log Ka] (Eq. 3) By taking the negative of the common logarithm of Equation 1, a useful relationship between pKa and pH called the Henderson‐Hasselbach equation can be derived. pH  pK a  log
A  -
HA 
(Eq. 4) If a solution of a weak acid is titrated to its endpoint, the anion, A–, species is present in the form of a salt. If this solution and the acid are mixed to give a 50:50 mixture then [HA] = [A–]. A solution that contains the weak acid, HA, and the anion, A–, of this weak acid from its salt is a buffer system. In a buffer solution in which [A–] = [HA], the measured pH of the solution equals to the pKa value. These conditions offer a convenient way to determine the pKa of an acid from pH measurements. Ka, is determined from pKa (Eq. 3). A. The Use of a pH Meter to Measure pH (See webpage for electrode calibration instructions) A pH meter gives more accurate measurements of pH than do indicators, but it uses more expensive equipment and must be calibrated, pH meters are generally similar in their operation Your instructor will provide some instructions for the use of the computer based pH system. A pH meter consists of a voltage measuring device and two electrodes. A reference electrode, usually a calomel electrode, provides a constant potential while glass electrode generates a potential that is proportional to the pH of the solution. When the two electrodes are placed in a solution, the voltage is displayed on a meter that is calibrated in pH units. The pH meter is calibrated against a standard solution of known pH before use. Care must be used in handing the electrodes. The thin membrane that is permeable to H3O+ ions is very fragile! In this procedure, a weak monoprotic acid solution will be titrated with a standardized solution of a strong base, NaOH. The pH will be recorded after the addition of small increments of base until the equivalence point is passed. A plot of pH on the vertical axis versus volume of NaOH on the horizontal axis gives a typical weak acid‐strong base titration curve (see text reference). At the point on the curve where the amount of base added is equal to one‐half the amount of base needed to reach the equivalence point (end point), the concentrations of the remaining acid, HA, and its anion, A–, are equal, and so the pH is equal to pKa. The concentration of the acid, MA, is calculated from the number of milliliters of the acid, VA, the concentration of the standard base, MB, and the number of milliliter of standard base, VB. MAVA = MBVB (True only for monoprotic acids and bases.) PROCEDURE B. Determination of Ka Values for Acetic Acid Using a pH Meter Step 1. Standardize the meter by following the procedures outlined in the laboratory techniques section of this lab manual. Step 2. Pipet 10.0 mL of the acetic acid solution into a 250 mL beaker and add 75 mL of distilled water and 2 drops of phenolphthalein indicator. Record this volume. (NOTE: Observe the volume of NaOH and the pH value when the indicator color changes for comparison with the pH curve you will plot.) Step 3. Immerse the electrodes in the solution and type the volume/pH measurement Step 4. Fill the buret nearly full with standard (0.100±0.010) M NaOH solution. Record the concentration of the NaOH solution, the initial buret reading and the initial pH of the acid solution in your notebook. Step 5. Add the NaOH solution in 1.0 mL increments and record the pH and the total volume of NaOH after each addition. Continue repeating this step until you have passed the equivalence point. (NOTE: Use the first sample to determine the approximate volume of standard NaOH solution required to reach the equivalence point.) Step 6. Plot your data as pH versus volume of NaOH solution added. Step 7. Repeat this titration once using new samples of the acid (i.e., repeat Steps 4‐5). Record the pH after each addition of NaOH. (a) Begin adding the NaOH solution in 2.0 mL increments until you have added about half of the required NaOH. (b) Then add the NaOH in 1.0 mL increment until you are within 2.0 mL of the equivalence point. (c) Add the NaOH in 0.10 mL increments until you are 1.0 mL beyond the equivalence point. (d) Finally add four 0.50 mL increments of NaOH solution. (e) Add 2 mL of NaOH 4 times to the solution to complete the titration curve. C. Determination of Ka Values for an Unknown Monoprotic Acid Using a pH Meter Step 8. Repeat Part B Steps 2 through 7 using one of the unknown weak monoprotic acid solutions as you did with acetic acid. D. Data Analysis Step 9. Plot your data as pH (vertical axis) versus mL of NaOH solution added Step 10. From your graphs, determine the exact volume of standard NaOH solution required to reach the equivalence point and then calculate the concentration of the acetic acid in your original solution. (Remember: MAVA = MBVB) Step 11. From your graphs, determine the value of Ka for CH3COOH and for the unknown acid. (Remember: the pKa = pH at the half‐way point of the titration.)