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pH Measurements, Equivalent Mass, and Ka Experiment Nine

pH Measurements and Determination of the Equivalent Mass and Ka of an Unknown Weak Acid

One of the more important properties of an aqueous solution is its concentration of hydrogen ion. The H+ or H3O+ ion has great effect on the solubility of many inorganic and organic species, on the nature of complex metallic cations found in solutions, and on the rates of many chemical reactions. It is important that we know how to measure the concentration of hydrogen ion and understand its effect on solution properties. For convenience, the concentration of H+ ion is frequently expressed as the pH of the solution rather than as molarity. The pH of a solution is defined by the following equation: pH = -log [H+] (1)

where the logarithm is taken to the base 10. For example, if [H+] is 1 x 10-4 moles per liter, the pH of the solution is 4. If the [H+] is 5 x 10-2 M, the pH is 1.3. Basic solutions can also be described in terms of pH. In aqueous solutions, the following equilibrium relationship exists: [H+] [OH-] = Kw = 1 x 10-14 at 25 ºC (2)

In pure water, [H+] = [OH-], so by equation 2, [H+] = 1 x 10-7 M. Therefore, the pH of pure water is ideally 7. Solutions in which [H+] > [OH-] are said to be acidic and will have a pH < 7; if [H+] < [OH-], the solution is basic and its pH > 7. In the first part of this experiment, you will determine the approximate pH of several solutions by using colorful acid-base indicators and comparing to various buffer solutions. The second part of the experiment involves predicting the pH of several salt solutions and then verifying your predictions using a pH meter. In the last part of the experiment, you will determine the gram equivalent mass (GEM) of an unknown acid, that is, the mass of the acid that supplies one mole of hydrogen ions. The acid, a solid crystalline substance, will be weighed out and titrated with a standard solution of sodium hydroxide. From the moles of base used and the mass of the acid, you will be able to determine the gram equivalent mass of the acid. Next, you will plot the titration curve of the acid, with pH on the vertical axis and the volume of NaOH on the horizontal axis. From this graph, you will be able to determine the value of the equilibrium constant for the dissociation of the acid, Ka. Acids are substances that contain ionizable hydrogen atoms within the molecule. Strong acids ionize completely, weak acids partially. The value of Ka, the equilibrium constant for the dissociation of the acid, is an indication of the strength of the acid. An acid may contain one or more ionizable hydrogen atoms in the molecule. The gram equivalent mass of an acid can be calculated from the molecular mass divided by the number of

Chemistry M01B Laboratory Manual pp. 48

pH Measurements, Equivalent Mass, and Ka Experiment Nine ionizable hydrogen atoms in a molecule. For example, hydrochloric acid, HCl, contains one ionizable hydrogen atom; the molecular mass is 36.45 g/mole, and the equivalent mass is also 36.45 g/mole. Sulfuric acid, H2SO4, contains two ionizable hydrogen atoms; the molecular mass is 98.07 g/mole, yet the equivalent mass is 49.04 g/mole. Thus, 36.45 g of HCl or 49.04 g of H2SO4 would provide you with one mole of H+ ions. The equivalent mass may be determined by titrating an acid with a standard solution of NaOH. Since one mole of NaOH will react with one mole of hydrogen ion, at the equivalence point, the following relation holds: Vb x Mb = moles base = moles H+ GEMa = grams acid/moles H+ where Vb is the volume of base, Mb is the molarity of base, grams acid is the mass of acid used, and GEMa is the gram equivalent mass of the acid. The concentration of the NaOH solution must be accurately known. To "standardize" the NaOH (that is, to find its exact molarity so it becomes a secondary standard), the NaOH is titrated against a solid acid, potassium hydrogen phthalate, sometimes abbreviated KHP (shown below). This acid is chosen because it possesses qualities of a primary standard which include a relatively large molar mass, high purity, unreactive with the

O OH OK O or KHC 8 H 4 O 4

atmosphere, one invariable reaction, and soluble in the chosen solvent. Other advantages of using KHP include its affordability, and it is relatively nontoxic compared to other possible choices. Sodium hydroxide cannot be used as a primary solid because it reacts with the atmosphere so it does not remain pure, and it has a relatively low molecular weight. The titration is thus followed using phenolphthalein as an indicator. A graph of pH versus mL of NaOH added can be drawn by carefully following the titration with a pH meter. There should be a significant change in pH in the vicinity of the equivalence point. Note that the equivalence point will probably NOT be at pH 7, but will be on the basic side (i.e. we are titrating a weak acid with a strong base). The value of the equilibrium constant for the dissociation of the acid can be obtained from the graph below (Figure 1). If we represent the dissociation of the acid as HA(aq) + H2O(l) H3O+(aq) + A-(aq), then the equilibrium expression is: Ka = [H3O+][A-]/[HA]. When the acid is HALF neutralized, [HA] = [A-], so these terms cancel in the above equation, and Ka = [H3O+]. Therefore, when the acid is half-neutralized, the pH = pKa.

Chemistry M01B Laboratory Manual pp. 49

pH Measurements, Equivalent Mass, and Ka Experiment Nine The point where pH is equal to pKa can be found from the graph. Refer to the following Figure 1.

pH Change During Titration

pH C A B Volume of NaOH

Figure 1. Titration of a Monoprotic Weak Acid with Sodium Hydroxide In this graph, A = Volume NaOH at equivalence point; B = ½ volume of A or the volume when half-neutralized; and C = pH when half-neutralized, or pKa. Procedure Part 1: pH of Acetic Acid Solutions, Unknown, and Buffer Solutions 1. Obtain two 3 x 4 spot plates. In the first spot plate, place 5 drops of 1.0 M, 0.10 M, and 0.010 M acetic acid in 4 consecutive wells, filling all 12 spots (see Figure 2).

Indicator 1 2 3 4

1.0 M

0.1 M

0.010 M

Figure 2. Schematic of first spot plate solution distribution 2. Choose one appropriate indicator and add one drop to one row of each concentration. Use a different indicator for each of the other three rows and repeat with one drop of indicator for each concentration. Estimate the pH of each concentration using the indicator bottle information.

Chemistry M01B Laboratory Manual pp. 50

pH Measurements, Equivalent Mass, and Ka Experiment Nine 3. Now use the second spot plate and place 5 drops of buffer solutions in four consecutive wells that are close to the estimated pH. Using the same indicators as before, match the colors with each acetic acid solution, and predict the pH to the nearest 0.5 unit. Finally, wash out the first spot plate containing the acetic acid concentrations. Add 5 drops of your unknown solution in 4 consecutive wells. Add one drop of an indicator to each well. Estimate the pH of your unknown to the nearest 0.5 unit by comparing the indicator colors of your unknown with colors of the same indicator in a buffer of known pH.

4.

Part 2: pH of Salt Solutions 1. 2. For the six salt solutions in the Data Sheet, estimate the pH as acidic, neutral, or basic. Record your predictions before proceeding to #2! Once your predictions are complete, read the pH meters immersed in the salt solutions and record the actual pH values on your data sheet. How closely do your predictions correlate with the actual experimental results? Make corrections if needed. Write the molecular equation, complete ionic equation, and net ionic equation for each of the salt solutions in the Data Sheet.

Part 3: GEM and Ka Determination of Unknown Weak Acid 1. Standardization of Approximately 0.10 M NaOH Since solid NaOH rapidly absorbs both H2O and CO2, a solution of exact molarity cannot be prepared by weighing the solid and diluting to volume. Instead, you must find its exact molarity by titrating it against a primary standard such as KHP. A. Accurately weigh about 0.5 g of KHP into a clean, dry 125 mL Erlenmeyer flask. Add about 30 mL of D.I. water and swirl until dissolved. Fill the buret with your NaOH solution. Open the stopcock briefly and shake the buret vertically to remove air bubbles from the tip. Add three drops of phenolphthalein solution to the acid in the flask and then titrate with the NaOH until the first trace of pink color persists for 30 seconds. Remember to constantly and gently swirl the flask. Look carefully and verify that all the KHP solid is dissolved. Rinse the sides of the flask with D.I. water and record the volume of NaOH used, estimating to the nearest 0.1 mL. Repeat two more times. If you use slightly more acid each time, the second and third titrations can be much more rapid than the first because you will know how much NaOH you can safely add before you get close to the end point. Calculate the average molarity of your NaOH solution.

B. C.

D.

E.

Chemistry M01B Laboratory Manual pp. 51

pH Measurements, Equivalent Mass, and Ka Experiment Nine 2. Determination of the Equivalent Mass of an Unknown Acid A. Accurately weigh a sample of your unknown acid (you will be given an unknown with a number and appropriate mass) using the sensitive balance. MAKE SURE TO RECORD WHICH UNKNOWN YOU ARE GIVEN. Dissolve in 30 mL D.I. water and titrate to the phenolphthalein end point as above. Repeat one more time. Choose a mass for the second sample so that the volume of NaOH needed will be about 40 mL if you are using a 50-mL buret. Calculate the gram equivalent mass of your sample.

B. C.

D. 3.

Determination of the Ka of an Unknown Acid A. B. C. D. On the analytical balance, weigh a sample of your acid that will require approximately 40 mL if you are using a 50-mL buret. Dissolve the sample in approximately 100 mL D.I. water in a 250-mL beaker. Set up a pH meter and electrode. Rinse the electrode well with D.I. water. NEVER wipe the electrode with a paper towel; this can damage the probe. Set the beaker on a magnetic stirrer. Clamp the pH electrode so it is submerged in your acid solution. Be sure the magnet does NOT hit the electrode. Titrate with standard base, recording the volume of base and pH of the solution during the titration. You should record the volume and pH every mL during the initial part of the titration, but as you get closer to an equivalence point, you must use smaller increments, approximately one drop in the vicinity of the equivalence point. Continue the titration curve at least 3 mL beyond the equivalence point. Graph your data, with pH on the vertical axis and volume NaOH on the horizontal axis. Make your graph large enough to reflect the care you took with the measurements of the pH and volume NaOH. From the graph, determine the pKa of the acid, that is, the pH where the acid is halfneutralized. Calculate the Ka value of your acid. Determine the volume of NaOH needed to reach the equivalence point. Use this value and each of the two values you recorded for the titration of your acid to determine the gram equivalent mass of your acid. Average the three gram equivalent mass values.

E.

F. G.

Chemistry M01B Laboratory Manual pp. 52

pH Measurements, Equivalent Mass, and Ka Experiment Nine Data and Questions Part 1: pH of Acetic Acid Solutions, Unknown, and Buffer Solutions 1. Enter in the appropriate space the name of the indicator used, the observed color for the listed indicator, and the estimated pH value from the pH paper for the acetic acid solutions and unknown (Note: HAc = HC2H3O2, acetic acid). Unknown #: _______ Indicator Used ____________ ____________ ____________ estimated pH 2. 1.0 M HAc(aq) ____________ ____________ ____________ ____________ 0.10 M HAc(aq) _____________ _____________ _____________ _____________ 0.010 M HAc(aq) ______________ ______________ ______________ ______________ Unknown _________ _________ _________ _________

Calculate the percent dissociation of the acetic acid solutions. SHOW YOUR CALCULATIONS! 1.0 M HAc(aq):

0.10 M HAc(aq):

0.010 M HAc(aq):

Part 2: pH of Salt Solutions 3. PREDICT whether each of the salt solutions below is expected to be acidic, neutral, or basic: NaCl _________ NaC2H3O2 _________ KNO3 ____________ Na2CO3 ___________ ZnCl2 _____________

NH4Cl ________

Chemistry M01B Laboratory Manual pp. 53

pH Measurements, Equivalent Mass, and Ka Experiment Nine 4. Using the pH meter immersed in each salt solution, determine the actual pH: NaCl _________ NaC2H3O2 _________ KNO3 ____________ Na2CO3 ___________ ZnCl2 _____________

NH4Cl ________ 5.

Write balanced MOLECULAR, IONIC, and NET-IONIC equations for the hydrolysis reactions of each salt solution. A. NaCl(aq): Molecular:

Ionic:

Net-Ionic:

B.

NaC2H3O2(aq) Molecular:

Ionic:

Net-Ionic:

C.

Na2CO3(aq) Molecular:

Ionic:

Net-Ionic:

Chemistry M01B Laboratory Manual pp. 54

pH Measurements, Equivalent Mass, and Ka Experiment Nine D. NH4Cl(aq) Molecular:

Ionic:

Net-Ionic:

E.

KNO3(aq) Molecular:

Ionic:

Net-Ionic:

F.

ZnCl2(aq) Molecular:

Ionic:

Net-Ionic:

Part 3: GEM and Ka Determination of Unknown Weak Acid 6. Prepare a pH profile graph of pH versus volume NaOH. Make certain to include this graph and the recorded data when submitting this report. What is the required volume to reach the equivalence point?

7.

8.

Calculate the pH at the half-equivalence point.

Chemistry M01B Laboratory Manual pp. 55

pH Measurements, Equivalent Mass, and Ka Experiment Nine 9. Calculate the pKa, Ka, and gram equivalent mass of your unknown acid. Unknown #: _______

10.

Why is the equivalence point NOT at pH 7?

Chemistry M01B Laboratory Manual pp. 56

pH Measurements, Equivalent Mass, and Ka Experiment Nine Advance Study Assignment: Equivalent Mass and Unknown Acid Determination 1. What is the equivalent mass of each of the following acids? B. KHCO3 C. H2SO3 A. HC2H3O2

2.

Calculate the molarity of a solution of sodium hydroxide if 23.6 mL is needed to neutralize 0.563 g of KHP.

3.

It is found that 24.6 mL of 0.116 M NaOH is needed to titrate 0.293 g of an unknown acid to the phenolphthalein end point. Calculate the equivalent mass of the acid.

4.

The following values were experimentally determined for the titration of 0.145 g of a weak acid with 0.100 M NaOH: Volume NaOH, mL 0.0 5.0 10.0 12.5 15.0 20.0 24.0 24.9 25.0 26.0 30.0 pH 2.88 4.15 4.58 4.76 4.93 5.36 6.14 7.15 8.73 11.29 11.96

Chemistry M01B Laboratory Manual pp. 57

pH Measurements, Equivalent Mass, and Ka Experiment Nine A. B. C. D. E. Graph the data. What is the required volume to reach the equivalence point? What is the pH at the half-equivalence point? Give the Ka and pKa value of the acid. Calculate the gram equivalent mass of the acid.

5.

The following acid-base indicators are available to follow the titration. Which of them would be most appropriate? Explain. Indicator Bromphenol blue Bromthymol blue Thymol blue Color Change Acid Form Base Form yellow blue yellow blue yellow blue pH Transition 3.0-5.0 6.0-7.6 8.0-9.6

Chemistry M01B Laboratory Manual pp. 58

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