014 SUSB Identification of an Unknown Weak Acid by pH Titration prepared by R. C. Kerber, J. W. Lauher, and M.J. Akhtar, SUNY at Stony Brook (Rev 10/08, RFS) Purpose of this Exercise: To identify an unknown weak acid by measuring its titration curve. From the titration curve, to determine the pKa and molar mass of the acid, which will allow identification by comparison with a list of common weak acids. Background Information Weak acids are all around us. Earlier in this course, we determined their presence in a variety of juices and beverages. Organic weak acids are formed in many bacterial decomposition reactions of fats and oils; for example, rancid butter smells of butyric acid and unwashed sweatsocks smell of isobutyric acid. Citrus juices derive much of their tangy flavor from citric acid, and tartaric acid may form beautiful crystals on the cork of a wine bottle. Boric acid is used in eyewash, and sodium hydrogen sulfate in toilet cleaner. Review: When a weak acid HA is dissolved in water, only a fraction of the acid molecules dissociate into H+ (aq) and A- (aq) ions. Most of the acid molecules remain intact. However, the molecules and ions do not remain fixed as such; over any given time interval, some molecules will dissociate and some ions will recombine. This occurrence of forward and reverse reactions leads to a condition of balance, called equilibrium. In an equilibrium situation, the net concentrations of the species involved do not change. The quantitative measure of the position of the equilibrium is called the equilibrium constant. For an acid dissociation, the equilibrium constant is called the acid dissociation constant, or Ka. It is defined as follows: Chemical Equation: Equilibrium Expression: HA H+(aq) + A-(aq) Ka = [ H+ ][ A- ] / [ HA ]
In the equilibrium expression, the use of square brackets indicates the concentration of the chemical entity inside the brackets, expressed in moles per liter (or millimoles per milliliter), and symbolized as M. So [ H+ ] means "concentration of H+ in moles per liter." In the case of a weak acid, the actual value of Ka is a small number, indicating that only a small fraction of the acid is dissociated into ions at any given time. For acetic acid, for example, Ka is about 1.75 x 10-5 in water at room temperature. Since use of exponential numbers is rather clumsy, we often use the negative log of the Ka, or pKa, instead of the Ka. The pKa of acetic acid is 4.75. How do we measure such a small number? A weak acid, even though it does not spontaneously dissociate to a very large extent, may react completely with a strong base such as hydroxide ion. You already know that you can determine the concentration of a weak acid by titrating with
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sodium hydroxide solution. When the number of moles of hydroxide added equals the number of moles of acid present, the pH (pH = -log [ H+ ]) rises abruptly, which can be detected by use of an indicator whose color changes with pH, or by use of a pH meter. For a sufficiently weak acid, what is the result of adding exactly half as many moles of hydroxide ion as there are moles of acid present? We might expect that half of the acid will be neutralized. The concentrations of HA and A- in the resulting half-neutralized solution will be equal: [ HA ] = [ A- ]. Let us call this concentration X, so that [ HA ] = [ A- ] = X. Then, going back to the definition of Ka, Ka = [ H+ ] [ A-] / [HA] = [ H+ ] X / X = [ H+ ]. This result shows that the Ka is equal to [ H+ ], the hydrogen ion concentration, in a solution where the concentrations of un-ionized acid, HA, and acid anion, A-, are equal. Also, pH = pKa. Therefore, to determine the Ka of an acid, all we need to do is to titrate the acid with NaOH solution, and to carefully record the pH as a function of volume of base added. We can plot the resulting data in the form of a graph, called a titration curve. When we reach the point of equivalence and the pH rises sharply, we know that the number of moles of base added equals the number of moles of acid originally present. If we look back at the titration curve to the point where we have added exactly half as much base as was required for complete neutralization, then the [ H+ ] at that point equals the Ka, and the pH at that point equals the pKa.*
TITRATION OF WEAK ACIDS
14 Half Titration Point 12 Equivalence Point
10
pKa = 8
8
pH
pKa = 6
6
pKa = 4
4
pKa = 2
2
Strong Acid
0
0.0 2.5 5.0 7.5 10.0 12.5 15.0 17.5 20.0 22.5 25.0 27.5 30.0 32.5 35.0 37.5
mL of BASE ADDED
A full discussion of this approximation is given on the web at http://www.ic.sunysb.edu/Class/che134/lectures/phpkahalftitr.pdf 34
Sample titration curves for several weak acids are shown above. Suppose 0.4606 g ( = 460.6 mg) of unknown acid was dissolved in water and titrated with 0.214 M. NaOH solution. The end point or equivalence point was reached when 25.00 mL of NaOH solution had been added to the acid solution. From this, we can calculate the molar mass (per acidic group) of the unknown acid: 25.00 mL x 0.214 mmols/mL = 5.350 mmol of OH5.350 mmol OH- 5.350 mmol acid molar mass of acid = 460.6 mg / 5.350 mmol = 86.1 Now looking at the overall shape of the titration curves, we note that at the beginning, when we have only weak acid in the solution, the pH is relatively low. As we add base, it rises. For very weak acids, the initial rise is sharp. There is a region half-way to the equivalence point when the pH changes only slightly upon addition of base. This is called the buffer region. In the exact middle, halfway to the equivalence point, corresponding to 12.50 mL (one-half of 25.00) of base added, we can read the pH from the smooth curve. If, in the example above, it turns out to be 4.8, that value is also the pKa of the weak acid in question. We now have two numerical values relating to the unknown acid, derived from a single experiment. Consulting the list of possible acids given in Table 1 at the end of this exercise, we can identify the unknown acid with a substantial degree of certainty. It is recommended that before coming to the laboratory, you prepare a table in your laboratory notebook for the data collection in the pH titration. To be sure that you plot pH against the appropriate (net) volume, you may wish to use a three-column table such as the one that follows. Buret Reading Net Volume Added pH Buret Reading Net Volume Added pH
Procedures I. Titration Curve CAUTION Sodium hydroxide solutions are very caustic and can cause prolonged and painful injuries to skin or eyes. Avoid all contact with these solutions. If accidents occur, flood the affected area with water for several minutes. Wear safety glasses at all times while in the laboratory.
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1. Review the procedures for use of pH meters in SUPL-006 2. Record the sample number of your unknown acid, and transfer the identifying sticker to your data sheet. Accurately weigh out approximately 0.2 g. of unknown into a clean Erlenmeyer flask by difference. Record the mass in your notebook. 3. Dissolve the acid sample in about 40 mL of distilled water, warming gently, if necessary, to insure complete dissolution. Cool the solution to room temperature. 4. Bring no more than 100 mL of standardized NaOH solution to your workplace. Keep your solution covered when not in use. Rinse a buret with about 10 mL of standardized NaOH solution, and then fill it to near the top of the graduated portion. Record the concentration of the base from the dispensing container, record and your initial buret reading. 5. In a preliminary run, add two or three drops of phenolphthalein solution to the flask, and titrate the acid solution to the point where a faint pink phenolphthalein color just persists. Calculate a molar mass based on this preliminary titration. 6. Calculate the weight range of your unknown that will require 25 + 3 mL of base. Weigh out, again in a clean 150 mL beaker, the appropriate amount of unknown acid, and record the mass in your notebook. After you have weighed the sample, compute the volume of base you should expect to use for this (actual) mass of acid. Keep this volume in mind as you perform the pH titration. 7. Again, dissolve the unknown acid in about 40 mL of distilled water, warming if necessary. 8. With the pH meter set at STANDBY, rinse the pH meter electrode with distilled water, using a squeeze bottle containing distilled water and a spare beaker to catch the washings. Place the electrode into your acid solution, so that the tip of the electrode is fully immersed, but not touching the bottom of the beaker. Refill your buret with standard NaOH solution, and mount it so that it will deliver solution into your beaker. Take the initial buret reading and record it. As you proceed, enter the results in the table in your notebook showing buret reading, total net volume added, and pH after each addition. 9. Turn the knob on the pH meter to the pH position, and record the initial pH reading. 10. Add a portion, perhaps 2-3 mL, of NaOH solution, and record the volume reading (as always read to the nearest 0.02 mL) and pH reading. Be sure to stir the solution with a stirring rod after each addition of base but before making the pH reading. 11. Add additional portions of NaOH solution, stir the solution and take volume and pH readings after each addition. The volume readings you record should be the actual buret readings, not the incremental ones. You will subtract the initial reading later.
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12. When you are within 2 mL of the equivalence point (refer to your calculation from part 6 above), the pH will rise more rapidly, and you should now add smaller increments of NaOH solution between readings. Within 1 mL of the equivalence point, and only in that region, you should take readings drop by drop. 13. After you have passed the equivalence point, add at least two or three additional portions of NaOH solution and continue to take readings as before, in order to define the upper region of the titration curve. Go at least 10 mL past the equivalence point. 14. Return the pH meter to STANDBY and remove the electrode from your solution. Rinse it with distilled water, and leave it standing in distilled water for the next user. 15. Using careful graphing technique (as described in SUPL-004#), construct a titration curve, remembering to subtract the initial buret reading from each subsequent reading to get actual volume added. Connect your experimental points with a smooth curve. If you choose to plot your titration curve using a computer, be sure that the graph you produce satisfies the guidelines for graphs described in SUPL-004 in particular, that it fills at least 2/3 of an 8 by 11 inch sheet of paper in landscape mode. 16. Determine from the graph the equivalence point, which is the point of steepest ascent of the curve. Record the volume of NaOH solution added at that point, and calculate the molar mass of your unknown acid from that volume. 17. Divide this volume by 2, and mark this half-neutralization point on your volume axis. Draw a vertical line which intercepts your titration curve, then a horizontal line from the point of intersection over to the pH axis. Determine the pH at this point; this is also the pKa of your unknown acid. 18. If the curve you obtain has significant irregularities, you probably did not stir the solution sufficiently after each addition of base. Repeat the pH titration. 19. Consult Table I to find an acid which matches yours in both molar mass and in pKa. A reasonable match in both parameters will constitute good evidence of identity. You should use your most reliable values of the equivalent molar mass and pKa in reaching your conclusion. You may wish to repeat the phenolphthalein titration if the equivalent molar mass computed from the first run differs substantially from that determined using the pH titration graph.
# SUPL-004
is available on the course web site.
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TABLE I: WEAK ACIDS
COMPOUND Acetic acid Propanoic acid Crotonic acid dl-Lactic acid Chloroacetic acid Maleic acid Succinic acid
FORMULA
MOLAR MASS 60.0 74.1 86.1 90.1 94.5 116.1 118.1 126.1 128.5 132.1 134.1 136.1 136.2 137.4 138.0 146.1 152.1 168.1 173.2 188.1 190.1 191.2 204.2
pKa# 4.8 4.9 4.7 3.9 2.9 1.8, 5.9 4.2, 5.6 1.2, 4.2 4.2 4.3, 5.2 3.4, 5.0 7.2, 12.7 1.9 1.9 7.2, 12.7 4.3, 4.4 3.4 3.0, 4.2 3.2 4.3 4.3 3.2 5.5
Oxalic acid dihydrate Potassium hydrogen oxalate Glutaric acid dl-Malic acid Potassium dihydrogenphosphate Potassium bisulfate Sodium bisulfate hydrate Sodium dihydrogenphosphate Adipic acid dl-Mandelic acid dl-Tartaric acid Sulfanilic acid Potassium hydrogen tartrate Sodium hydrogen tartrate Sulfanilic acid hydrate Potassium hydrogen phthalate
HC2H3O2 HC3H5O2 HC4H5O2 HC3H5O3 HC2ClH2O2 H2C4H2O4 H2C4H4O4 H2C2O4.2H2O KHC2O4 H2C5H6O4 HC4H4O5 KH2PO4 KHSO4 NaHSO4.H2O NaH2PO4.H2O H2C6H8O4 HC8H7O3 H2C4H4O6 HC6H6NO3S KHC4H4O6 NaHC4H4O6.H2O HC6H6NO3S.H2O KHC8O8
# The pKa values in the table should be viewed as approximate. Measured values may differ from the tabulated values by 0.5 units or more.
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Notebook Grade: _____ SUSB-014 Data Sheet Safety Grade: _____
____________________________________________________________________________ Name Section Date Conc. of NaOH solution from label _________M. Run 1 _________g _________g _________g _________mg _________M. Run 2 (pH titration) _________g _________g _________g _________mg _________mL _________mL _________mL _________mL _________mL(graph)
Initial Mass of vial Final Mass of vial Mass of acid sample Mass of acid sample Calculated volume of NaOH required Final buret reading Initial buret reading Net Volume of NaOH sol'n. at equivalence point
Millimoles of NaOH Millimoles of acid Molar mass of unknown acid
_________mmol _________mmol
_________mmol _________mmol
_________mg/mmol _________mg/mmol
Weight of acid for 25 mL of NaOH in Part 2
_________mg
Volume of NaOH sol'n at half-equivalence point pH of solution at half-equivalence point pKa of unknown acid
_________mL(graph) _________(graph) _________
Identity of unknown acid:
Attach your pH titration curve.
Be sure the relevant points are clearly labeled.
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SUSB-014 Pre-Laboratory Assignment ________________________________________ ________________ ____________________ Name Section Date 1. Adipic acid has a pKa of 4.40. What is its Ka?
2. Sketch the titration curve for titrating 0.1461 g. of dl-mandelic acid with 0.05 M. sodium hydroxide solution. Use data from Table 1. Mark and label the end point and the half titration point.
3. An unknown acid (0.2922 g) was titrated with 0.100 M. sodium hydroxide. The equivalence point required 20.00 mL. The pH at the point where 10.00 mL had been added was 4.40. Identify the unknown acid, using data from Table 1.
4. X mg of an unknown acid was titrated with 0.110 M sodium hydroxide. Calculate X by subtracting the last two digits of you student ID in mg from 400 mg. The equivalence point is at 25.00 mL. What is the molar mass of the unknown acid?
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