Equilibrium: Titration of an Amino Acid

• To acquire skill in volumetric titration
• To explore the acid-base behavior of an amino acid.

According to the Bronsted view, an acid is a substance capable of donating a proton, H⁺, to a base. Thus an acid is a proton donor, and a base is a proton acceptor. Reaction between a generic acid, HA (note that we represent the donatable proton explicitly) and a generic base, B, is shown in equation 1:

B uses a non-bonding pair of electrons to attach the proton; this pair is shown in the equation. The reverse of reaction 1 also involves proton transfer, with BH⁺ serving as the acid and A^- as the base. The pair of species, HA and A- are related by gain or loss of H⁺. They are therefore called a conjugate acid-base pair. B and BH⁺ are also a conjugate acid-base pair.

Our primary interest in acids in this experiment involves their reaction with water, shown in equation 2.

In this reaction, water functions as a Bronsted base. The extent of reaction 2 is indicated quantitatively using the equilibrium constant, K_eq. The equilibrium constant expression for 2 is given in 3:

The equilibrium constant for reaction of an acid with water is usually symbolized K_a, to remind us of the type of reaction being dealt with. The reactant water, since it is present in huge concentration and is thus essentially a pure liquid, is not included in the K_a expression. The strength of an acid in aqueous solution is defined in terms of the magnitude of K_a for reaction 2. Strong acids have K_a values larger than 1; weak acids have K_a values less than 1. The common strong acids are HCl (hydrochloric), HBr (hydrobromic), HI (hydroiodic), HNO₃ (nitric), H₂SO₄ (sulfuric), HClO₃ (chloric) and HClO₄ (perchloric). All other acids that we will commonly encounter are weak; i.e., their reaction with water according to 2 occurs to only a minor extent.

The equilibrium established when a weak acid reacts with water according to 2 can be explored using a procedure called titration, involving the following steps:

• A precisely known volume of acid solution is measured out. Such a precisely known volume is called an aliquot, and may be measured using a volumetric pipet, an item of glassware which, when filled with the solution of interest, delivers precisely the indicated volume of solution. Convenient pipet sizes for aliquot measurement range from 1.00 to 25.00 mL.
• To the aliquot of acid solution, a solution of base of precisely known concentration (called standard base) is added in such a way that the volume added can be precisely measured. Most often, the base is dispensed from an item of precision glassware called a buret. It is also possible to dispense base using an automatic pipettor, which can deliver accurately any selected volume within its range of manufacture. The process of adding base to acid is the actual titration. The substance added from the buret or pipettor (usually the base), is called the titrant. Addition of titrant is continued until the moles of OH^- added is precisely equal to the moles of H⁺ in the acid. The point of the titration where moles OH^- added = moles H⁺ in acid is called the equivalence point.
• From the volume of base added and its known concentration, we calculate the concentration of acid in the original solution.

Titration involves a very important practical problem: how is the equivalence point detected? Unless we can detect the equivalence point, we will not know when to stop adding titrant. Clearly, the pH of the solution must change as the titration proceeds. At the beginning of the process, before base is added, the pH of the solution is fairly low because it contains acid. As titration proceeds, acid is neutralized by the added base, and pH rises. Addition of base after all of the acid has been neutralized produces a basic solution, with a high pH. During the titration, then, pH runs the gamut from low to high. We can detect the equivalence point from the manner in which the pH change occurs.

Two methods are commonly used to detect the equivalence point in a titration. In the first, an appropriate acid-base indicator is used to signal the equivalence point via a color change. In the second, an instrument called a pH meter is used to monitor the pH of the solution as base is added during the titration. A pH meter responds to the electrical potential of an electrode immersed in the solution being titrated. This potential is a function of [H₃O⁺] in the solution. A plot of pH versus the volume of titrant added to the solution gives the so-called titration curve. The experimental curve for titration of 40.00 mL of 0.1000 M HCl with 0.1000 M NaOH is shown in Figure 1. We observe that

• The curve is shaped something like the letter "S". All titration curves have this characteristic shape.
• The total change in pH during the titration is about 12 units, but almost all of this pH change occurs over a very small range of base volume centered at 40 mL, which corresponds to the equivalence point. Thus the large pH change occurs right at the equivalence point! More exactly, the equivalence point corresponds with the inflection point of the curve, which is the point at which the slope of the curve is a maximum.

This provides us with the second method for determining the equivalence point: we successively add small volumes of base, measure pH after each addition, and plot the titration curve, from which we may find V_base at the inflection point (the equivalence point). Moles of acid in the original aliquot is calculated as follows:

In this experiment, we will use the pH meter method to explore the acid-base behavior of a number of amino acids. Amino acids are molecules that contain both a base site (an -NH₂ group) and an acid site (a -COOH group). The structure shown is generic for amino acids. Individual acids differ only in the identity of the group, -R.

When an amino acid is dissolved in water, the proton from the -COOH group transfers to the -NH₂ end of the molecule, because the NH₂ group is a stronger base than -COO-. The resulting structure is called a Zwitterion.

A titration of the Zwitterion with standard NaOH would provide the K_a value for the -NH₃⁺ acid, which would be expected to be similar to that of NH₄⁺ (pK_a = 9.25). However, it would be nice to also obtain the K_a value for the -COOH acid. It is possible to generate this acid in solution by adding strong acid to the Zwitterion. The strong acid transfers a proton to the -COO^- base of the Zwitterion, resulting in a cation.

Titration of a solution of this cation with standard NaOH should then yield two equivalence points, one for each acid. It should thus be possible to measure both desired K_a values.

1. If you were asked to estimate the value of pK_a for the -COOH end of an amino acid, what information would you require?
2. Why is it necessary to add exactly the stoichiometric amount of strong acid to your Zwitterion solution?
3. What is the significance of the pH at which the concentrations of conjugate acid and base are the same?
4. Why are the equivalence points less distinct than in the titration of, say, acetic acid with NaOH?
5. Do the measured pK_a values deviate from "expected" values? Can you explain why?

• Buret (10- or 25-mL)
• 1.000-mL automatic pipettor
• pH meter with pH electrode
• stirring rod or magnetic stir plate with stir bar
• 2 50-mL beakers
• Solutions:
  • 0.10 M KHP, standardized
  • ~0.1 M NaOH
  • ~0.1 M HCl
  • Amino acids
    • glycine
    • alanine

Safety goggles must be worn at all times in the laboratory. Dilute solutions of acids and bases are irritating to the skin. In the event of skin contact, rinse thoroughly with plenty of water.

Record all data in your notebook.

Your team will be assigned a pH meter and a particular amino acid to investigate.

Obtain about 30 mL of ~0.1 M NaOH in a clean dry beaker or Erlenmeyer. Be sure to label the container, and cover it with a watch glass. Obtain about 15 mL of standard KHP solution in another clean container, label it, and cover it. Record the exact concentration of the KHP solution. Keep the reservoir containers covered with watch glasses at all times to retard evaporation and exclude CO₂ from the air.

Obtain a buret, and inspect it to be sure that it is clean on the inside walls. Rinse it 3 times with 1-2 mL portions of ~0.1 M NaOH solution, then fill to near the 0.00 mark. Clear all air from the tip of the buret.

Pipet 5.0 mL of standard KHP solution into a clean dry Erlenmeyer, and add 2 drops of phenolphthalein indicator solution. Read the level of titrant in the buret by matching the lowest point of the meniscus with the graduation marks on the buret. You should read the volume level to the nearest 0.01 mL. Titrate with NaOH solution to the faintest possible pink endpoint, and again read the level of titrant in the buret. The volume of base delivered is the difference between your initial and final volume readings. Repeat the titration using 5.0-mL aliquots of standard KHP solution until you have three results that differ from one another by no more than 1%. (For example, values of 6.20, 6.26, and 6.23 mL have a range of 0.06 mL from smallest to largest. This is slightly less than 1% of the median value (0.06/6.23 * 100 = 0.96%, and so is just within acceptable limits of precision.) Using the average of your three volumes and the known volume and concentration of the KHP solution, calculate the molarity of the NaOH solution to 3 significant figures. The NaOH solution has now been standardized.

Obtain about 30 mL of ~0.1 M HCl in a clean dry beaker or Erlenmeyer. Label and cover the container. Pipet exactly 5.00 mL of the HCl solution into a clean Erlenmeyer flask, add 2 drops of phenolphthalein indicator solution, and titrate to the pink endpoint with standard NaOH. Repeat the standardization until you have three results with a range of less than 1% of the median value. Using the average of your three volumes, the known molarity of the NaOH solution (calculated above) and the known volume of HCl solution, calculate the molarity of the HCl solution to 3 significant figures. The HCl solution has now been standardized.

Calculate the mass of amino acid needed to prepare 50 mL of a 0.075 M solution of your assigned amino acid. Weigh reasonably close to this amount of amino acid into a 50-mL volumetric flask. NOTE: it is not important that you hit your target mass exactly; however, it is important that you know EXACTLY what mass of amino acid is in the flask. Calculate the number of moles of amino acid in the flask. Then calculate the volume of standardized HCl solution needed to provide exactly this number of moles of HCl. Add this volume of standardized HCl solution to the flask using a graduated pipet of appropriate capacity. Swirl the contents of the flask to dissolve the amino acid, then add deionized water to the mark on the flask. Homogenize the solution by shaking the heck out of it. Addition of HCl to the amino acid has now protonated the amino acid carboxylate group.

(NOTE: this step may or may not be necessary; check with the instructor.) Before beginning, calibrate the pH meter using standard buffer solutions of pH 4 and pH 7, or pH 7 and pH 10. Your instructor will provide instructions for doing the calibration and will tell you which two buffers to use.

Transfer exactly 10.00 mL of the amino acid solution to a clean 50-mL beaker, and add deionized water until the total volume is 25-30 mL. Place the beaker on the top plate of a magnetic stirrer. Place a 1-inch stir bar in the beaker. Rinse and dry the pH electrode and submerge it in the solution containing protonated amino acid. Make sure that the tip of the electrode is clear of the magnetic stir bar in the beaker, then start the stirrer. The rotation rate should be reasonably fast, but not so vigorous that splashing of the solution occurs. RECORD THE INITIAL pH OF THE SOLUTION. Initiate the pH titration by adding 0.500 mL of NaOH solution from an automatic pipettor that has been dialed to deliver exactly this volume. Make sure that the base enters the solution at a point away from the pH electrode. When the pH meter reading is stable--that is, does not fluctuate or steadily change in one direction--read and record the pH and the volume of NaOH solution added. Continue adding NaOH 0.50 mL at a time, recording pH and total volume of NaOH added after each addition, until the total added volume of base is about 80% of the amount required to titrate the -COOH proton. (IMPORTANT NOTE: Some amino acids have two -COOH groups, the normal alpha one, and another one in the R group. For these amino acids, twice as much NaOH will be needed to reach the first equivalence point as for most amino acids.) At this point, reset the pipettor to add the NaOH in smaller increments of first 0.20, then 0.10 mL. Continue adding NaOH incrementally until you are past the -COOH equivalence point. Go back to adding NaOH 0.5 mL at a time until you have added 80% of the amount required to titrate the -NH₃+ proton. At this point, add smaller increments as above until you are at least 2 mL past the second equivalence point. When finished, discard the solution. Rinse the pH electrode and beaker with distilled water, and dry both.

Repeat the titration at least once more, twice if time allows.

When finished, rinse the pH electrode with distilled water, dry it, cap it, and turn the meter off. Turn off the magnetic stirrer. Clean all glassware and put it in the drying oven.

Spent titration solutions may be flushed down the drain with plenty of water.

Clean-up. When you have finished all of your work:

• Clean all glassware with Alconox and water and shake dry.
• Rinse Pasteur pipets by inserting the nozzle of your wash bottle in the top end and squirting water through. Aspirate or oven dry.
• Return all borrowed equipment to the instructor before leaving lab.