Consider the titration of 80.0 mL of 0.100 M Ba(OH)2 by

What are the major species in solution after NaHSO4 is dissolved in water? What happens to the pH of the solution as more NaHSO4 is added? Why? Would the results vary if baking soda (NaHCO3) were used instead?

A friend asks the following: Consider a buffered solution made up of the weak acid HA and its salt NaA. If a strong base like NaOH is added, the HA reacts with the OH to form A. Thus the amount of acid (HA) is decreased, and the amount of base (A) is increased. Analogously, adding HCl to the buffered solution forms more of the acid (HA) by reacting with the base (A). Thus how can we claim that a buffered solution resists changes in the pH of the solution? How would you explain buffering to this friend?

Mixing together solutions of acetic acid and sodium hydroxide can make a buffered solution. Explain. How does the amount of each solution added change the effectiveness of the buffer?

Could a buffered solution be made by mixing aqueous solutions of HCl and NaOH? Explain. Why isnt a mixture of a strong acid and its conjugate base considered a buffered solution?

Sketch two pH curves, one for the titration of a weak acid with a strong base and one for a strong acid with a strong base. How are they similar? How are they different? Account for the similarities and the differences.

Sketch a pH curve for the titration of a weak acid (HA) with a strong base (NaOH). List the major species and explain how you would go about calculating the pH of the solution at various points, including the halfway point and the equivalence point

You have a solution of the weak acid HA and add some HCl to it. What are the major species in the solution? What do you need to know to calculate the pH of the solution, and how would you use this information? How does the pH of the solution of just the HA compare with that of the final mixture? Explain.

You have a solution of the weak acid HA and add some of the salt NaA to it. What are the major species in the solution?

The common ion effect for weak acids is to significantly decrease the dissociation of the acid in water. Explain the common ion effect.

Consider a buffer solution where [weak acid]  [conjugate base]. How is the pH of the solution related to the pKa value of the weak acid? If [conjugate base]  [weak acid], how is pH related to pKa?

A best buffer has about equal quantities of weak acid and conjugate base present as well as having a large concentration of each species present. Explain.

Consider the following four titrations. i. 100.0 mL of 0.10 M HCl titrated by 0.10 M NaOH ii. 100.0 mL of 0.10 M NaOH titrated by 0.10 M HCl iii. 100.0 mL of 0.10 M CH3NH2 titrated by 0.10 M HCl iv. 100.0 mL of 0.10 M HF titrated by 0.10 M NaOH Rank the titrations in order of: a. increasing volume of titrant added to reach the equivalence point. b. increasing pH initially before any titrant has been added. c. increasing pH at the halfway point in equivalence. d. increasing pH at the equivalence point. How would the rankings change if C5H5N replaced CH3NH2 and if HOC6H5 replaced HF?

Figure 15.4 shows the pH curves for the titrations of six different acids by NaOH. Make a similar plot for the titration of three different bases by 0.10 M HCl. Assume 50.0 mL of 0.20 M of the bases and assume the three bases are a strong base (KOH), a weak base with Kb 1  105 , and another weak base with Kb 1  1010.

Acidbase indicators mark the end point of titrations by magically turning a different color. Explain the magic behind acidbase indicators.

A certain buffer is made by dissolving NaHCO3 and Na2CO3 in some water. Write equations to show how this buffer neutralizes added H and OH.

A buffer is prepared by dissolving HONH2 and HONH3NO3 in some water. Write equations to show how this buffer neutralizes added H and OH.

Calculate the pH of each of the following solutions. a. 0.100 M propanoic acid (HC3H5O2, Ka 1.3  105 ) b. 0.100 M sodium propanoate (NaC3H5O2) c. pure H2O d. a mixture containing 0.100 M HC3H5O2 and 0.100 M NaC3H5O2

Calculate the pH of each of the following solutions. a. 0.100 M HONH2 (Kb 1.1  108 ) b. 0.100 M HONH3Cl c. pure H2O d. a mixture containing 0.100 M HONH2 and 0.100 M HONH3Cl

Compare the percent dissociation of the acid in Exercise 17a with the percent dissociation of the acid in Exercise 17d. Explain the large difference in percent dissociation of the acid.

Compare the percent ionization of the base in Exercise 18a with the percent ionization of the base in Exercise 18d. Explain any differences.

Calculate the pH after 0.020 mol HCl is added to 1.00 L of each of the four solutions in Exercise 17

Calculate the pH after 0.020 mol HCl is added to 1.00 L of each of the four solutions in Exercise 18.

Calculate the pH after 0.020 mol NaOH is added to 1.00 L of each of the four solutions in Exercise 17.

Calculate the pH after 0.020 mol NaOH is added to 1.00 L of each of the solutions in Exercise 18.

Which of the solutions in Exercise 17 shows the least change in pH upon the addition of acid or base? Explain.

Which of the solutions in Exercise 18 is a buffered solution?

Calculate the pH of a solution that is 1.00 M HNO2 and 1.00 M NaNO2.

Calculate the pH of a solution that is 0.60 M HF and 1.00 M KF

Calculate the pH after 0.10 mol NaOH is added to 1.00 L of the solution in Exercise 27, and calculate the pH after 0.20 mol HCl is added to 1.00 L of the solution in Exercise 27.

Calculate the pH after 0.10 mol NaOH is added to 1.00 L of the solution in Exercise 28, and calculate the pH after 0.20 mol HCl is added to 1.00 L of the solution in Exercise 28

Calculate the pH of each of the following buffered solutions. a. 0.10 M acetic acid/0.25 M sodium acetate b. 0.25 M acetic acid/0.10 M sodium acetate c. 0.080 M acetic acid/0.20 M sodium acetate d. 0.20 M acetic acid/0.080 M sodium acetate

Calculate the pH of each of the following buffered solutions. a. 0.50 M C2H5NH2/0.25 M C2H5NH3Cl b. 0.25 M C2H5NH2/0.50 M C2H5NH3Cl c. 0.50 M C2H5NH2/0.50 M C2H5NH3Cl

Calculate the pH of a buffer solution prepared by dissolving 21.5 g benzoic acid (HC7H5O2) and 37.7 g sodium benzoate in 200.0 mL of solution.

A buffered solution is made by adding 50.0 g NH4Cl to 1.00 L of a 0.75 M solution of NH3. Calculate the pH of the final solution. (Assume no volume change.)

Calculate the pH after 0.010 mol gaseous HCl is added to 250.0 mL of each of the following buffered solutions. a. 0.050 M NH3/0.15 M NH4Cl b. 0.50 M NH3/1.50 M NH4Cl Do the two original buffered solutions differ in their pH or their capacity? What advantage is there in having a buffer with a greater capacity?

An aqueous solution contains dissolved C6H5NH3Cl and C6H5NH2. The concentration of C6H5NH2 is 0.50 M and pH is 4.20. a. Calculate the concentration of C6H5NH3 in this buffer solution. b. Calculate the pH after 4.0 g NaOH(s) is added to 1.0 L of this solution. (Neglect any volume change.)

Calculate the mass of sodium acetate that must be added to 500.0 mL of 0.200 M acetic acid to form a pH 5.00 buffer solution.

What volumes of 0.50 M HNO2 and 0.50 M NaNO2 must be mixed to prepare 1.00 L of a solution buffered at pH 3.55?

Consider a solution that contains both C5H5N and C5H5NHNO3. Calculate the ratio [C5H5N][C5H5NH] if the solution has the following pH values. a. pH 4.50 c. pH 5.23 b. pH 5.00 d. pH 5.50 40.

Calculate the ratio [NH3][NH4 ] in ammonia/ammonium chloride buffered solutions with the following pH values: a. pH 9.00 c. pH 10.00 b. pH 8.80 d. pH 9.60 41.

Consider the acids in Table 14.2. Which acid would be the best choice for preparing a pH 7.00 buffer? Explain how to make 1.0 L of this buffer.

Consider the bases in Table 14.3. Which base would be the best choice for preparing a pH 5.00 buffer? Explain how to make 1.0 L of this buffer.

Calculate the pH of a solution that is 0.40 M H2NNH2 and 0.80 M H2NNH3NO3. In order for this buffer to have pH pKa, would you add HCl or NaOH? What quantity (moles) of which reagent would you add to 1.0 L of the original buffer so that the resulting solution has pH pKa?

Calculate the pH of a solution that is 0.20 M HOCl and 0.90 M KOCl. In order for this buffer to have pH pKa, would you add HCl or NaOH? What quantity (moles) of which reagent would you add to 1.0 L of the original buffer so that the resulting solution has pH pKa?

Which of the following mixtures would result in buffered solutions when 1.0 L of each of the two solutions are mixed? a. 0.1 M KOH and 0.1 M CH3NH3Cl b. 0.1 M KOH and 0.2 M CH3NH2 c. 0.2 M KOH and 0.1 M CH3NH3Cl d. 0.1 M KOH and 0.2 M CH3NH3Cl

Which of the following mixtures would result in a buffered solution when 1.0 L of each of the two solutions are mixed? a. 0.2 M HNO3 and 0.4 M NaNO3 b. 0.2 M HNO3 and 0.4 M HF c. 0.2 M HNO3 and 0.4 M NaF d. 0.2 M HNO3 and 0.4 M NaOH

What quantity (moles) of NaOH must be added to 1.0 L of 2.0 M HC2H3O2 to produce a solution buffered at each pH? a. pH pKa b. pH 4.00 c. pH 5.00 4

Calculate the number of moles of HCl(g) that must be added to 1.0 L of 1.0 M NaC2H3O2 to produce a solution buffered at each pH. a. pH pKa b. pH 4.20 c. pH 5.00 A

Consider the titration of a generic weak acid HA with a strong base that gives the following titration curve:On the curve, indicate the points that correspond to the following: a. the stoichiometric (equivalence) point b. the region with maximum buffering c. pH pKa d. pH depends only on [HA] e. pH depends only on [A] f. pH depends only on the amount of excess strong base added

Sketch the titration curve for the titration of a generic weak base B with a strong acid. The titration reaction is On this curve, indicate the points that correspond to the following: a. the stoichiometric (equivalence) point b. the region with maximum buffering c. pH pKa d. pH depends only on [B] e. pH depends only on [BH] f. pH depends only on the amount of excess strong acid added

Consider the titration of 40.0 mL of 0.200 M HClO4 by 0.100 M KOH. Calculate the pH of the resulting solution after the following volumes of KOH have been added. a. 0.0 mL d. 80.0 mL b. 10.0 mL e. 100.0 mL c. 40.0 mL

Consider the titration of 80.0 mL of 0.100 M Ba(OH)2 by 0.400 M HCl. Calculate the pH of the resulting solution after the following volumes of HCl have been added. a. 0.0 mL d. 40.0 mL b. 20.0 mL e. 80.0 mL c. 30.0 mL

Consider the titration of 100.0 mL of 0.200 M acetic acid (Ka 1.8  105 ) by 0.100 M KOH. Calculate the pH of the resulting solution after the following volumes of KOH have been added. a. 0.0 mL d. 150.0 mL b. 50.0 mL e. 200.0 mL c. 100.0 mL f. 250.0 mL

Consider the titration of 100.0 mL of 0.100 M H2NNH2 (Kb 3.0  106 ) by 0.200 M HNO3. Calculate the pH of the resulting solution after the following volumes of HNO3 have been added. a. 0.0 mL d. 40.0 mL b. 20.0 mL e. 50.0 mL c. 25.0 mL f. 100.0 mL

Lactic acid is a common by-product of cellular respiration and is often said to cause the burn associated with strenuous activity. A 25.0-mL sample of 0.100 M lactic acid (HC3H5O3, pKa 3.86) is titrated with 0.100 M NaOH solution. Calculate the pH after the addition of 0.0 mL, 4.0 mL, 8.0 mL, 12.5 mL, 20.0 mL, 24.0 mL, 24.5 mL, 24.9 mL, 25.0 mL, 25.1 mL, 26.0 mL, 28.0 mL, and 30.0 mL of the NaOH. Plot the results of your calculations as pH versus milliliters of NaOH added.

Repeat the procedure in Exercise 55, but for the titration of 25.0 mL of 0.100 M propanoic acid (HC3H5O2, Ka 1.3  105 ) with 0.100 M NaOH.

Repeat the procedure in Exercise 55, but for the titration of 25.0 mL of 0.100 M NH3 (Kb 1.8  105 ) with 0.100 M HCl.

Repeat the procedure in Exercise 55, but for the titration of 25.0 mL of 0.100 M pyridine with 0.100 M hydrochloric acid (Kb for pyridine is 1.7  109 ). Do not indicate the points at 24.9 and 25.1 mL.

Calculate the pH at the halfway point and at the equivalence point for each of the following titrations. a. 100.0 mL of 0.10 M HC7H5O2 (Ka 6.4  105 ) titrated by 0.10 M NaOH b. 100.0 mL of 0.10 M C2H5NH2 (Kb 5.6  104 ) titrated by 0.20 M HNO3 c. 100.0 mL of 0.50 M HCl titrated by 0.25 M NaOH

In the titration of 50.0 mL of 1.0 M methylamine, CH3NH2 (Kb 4.4  104 ), with 0.50 M HCl, calculate the pH under the following conditions. a. after 50.0 mL of 0.50 M HCl has been added b. at the stoichiometric point

You have 75.0 mL of 0.10 M HA. After adding 30.0 mL of 0.10 M NaOH, the pH is 5.50. What is the Ka value of HA?

A student dissolves 0.0100 mol of an unknown weak base in 100.0 mL water and titrates the solution with 0.100 M HNO3. After 40.0 mL of 0.100 M HNO3 was added, the pH of the resulting solution was 8.00. Calculate the Kb value for the weak base.

Two drops of indicator HIn (Ka 1.0  109 ), where HIn is yellow and In is blue, are placed in 100.0 mL of 0.10 M HCl. a. What color is the solution initially? b. The solution is titrated with 0.10 M NaOH. At what pH will the color change (yellow to greenish yellow) occur? c. What color will the solution be after 200.0 mL NaOH has been added?

Methyl red has the following structure:It undergoes a color change from red to yellow as a solution gets more basic. Calculate an approximate pH range for which methyl red is useful. What is the color change and the pH at the color change when a weak acid is titrated with a strong base using methyl red as an indicator? What is the color change and the pH at the color change when a weak base is titrated with a strong acid using methyl red as an indicator? For which of these two types of titrations is methyl red a possible indicator?

Potassium hydrogen phthalate, known as KHP (molar mass 204.22 g/mol), can be obtained in high purity and is used to determine the concentration of solutions of strong bases by the reaction If a typical titration experiment begins with approximately 0.5 g KHP and has a final volume of about 100 mL, what is an appropriate indicator to use? The pKa for HP is 5.51.

A certain indicator HIn has a pKa of 3.00 and a color change becomes visible when 7.00% of the indicator has been converted to In. At what pH is this color change visible?

Which of the indicators in Fig. 15.8 could be used for the titrations in Exercises 51 and 53?

Which of the indicators in Fig. 15.8 could be used for the titrations in Exercises 52 and 54?

Which of the indicators in Fig. 15.8 could be used for the titrations in Exercises 55 and 57?

Which of the indicators in Fig. 15.8 could be used for the titrations in Exercises 56 and 58?

Estimate the pH of a solution in which bromcresol green is blue and thymol blue is yellow. (See Fig. 15.8.)

Estimate the pH of a solution in which crystal violet is yellow and methyl orange is red. (See Fig. 15.8.)

A solution has a pH of 7.0. What would be the color of the solution if each of the following indicators were added? (See Fig. 15.8.) a. thymol blue c. methyl red b. bromthymol blue d. crystal violet

A solution has a pH of 4.5. What would be the color of the solution if each of the following indicators were added? (See Fig. 15.8.) a. methyl orange b. alizarin c. bromcresol green d. phenolphthalein

Tris(hydroxymethyl)aminomethane, commonly called TRIS or Trizma, is often used as a buffer in biochemical studies. Its buffering range is pH 7 to 9, and Kb is 1.19  106 for the aqueous reaction TRIS TRISH a. What is the optimal pH for TRIS buffers? b. Calculate the ratio [TRIS][TRISH] at pH 7.00 and at pH 9.00. c. A buffer is prepared by diluting 50.0 g TRIS base and 65.0 g TRIS hydrochloride (written as TRISHCl) to a total volume of 2.0 L. What is the pH of this buffer? What is the pH after 0.50 mL of 1

Phosphate buffers are important in regulating the pH of intracellular fluids at pH values generally between 7.1 and 7.2. a. What is the concentration ratio of H2PO4  to HPO4 2 in intracellular fluid at pH 7.15? b. Why is a buffer composed of H3PO4 and H2PO4  ineffective in buffering the pH of intracellular fluid?

Carbonate buffers are important in regulating the pH of blood at 7.40. If the carbonic acid concentration in a sample of blood is 0.0012 M, determine the bicarbonate ion concentration required to buffer the pH of blood at pH 7.40.

When a person exercises, muscle contractions produce lactic acid. Moderate increases in lactic acid can be handled by the blood buffers without decreasing the pH of blood. However, excessive amounts of lactic acid can overload the blood buffer system, resulting in a lowering of the blood pH. A condition called acidosis is diagnosed if the blood pH falls to 7.35 or lower. Assume the primary blood buffer system is the carbonate buffer system described in Exercise 77. Calculate what happens to the ratio in blood when the pH decreases from 7.40 to 7.35.

The active ingredient in aspirin is acetylsalicylic acid. A 2.51-g sample of acetylsalicylic acid required 27.36 mL of 0.5106 M NaOH for complete reaction. Addition of 13.68 mL of 0.5106 M HCl to the flask containing the aspirin and the sodium hydroxide produced a mixture with pH 3.48. Determine the molar mass of acetylsalicylic acid and its Ka value. State any assumptions you must make to reach your answer.

One method for determining the purity of aspirin (C9H8O4) is to hydrolyze it with NaOH solution and then to titrate the remaining NaOH. The reaction of aspirin with NaOH is as follows: Aspirin Salicylate ion Acetate ion A sample of aspirin with a mass of 1.427 g was boiled in 50.00 mL of 0.500 M NaOH. After the solution was cooled, it took 31.92 mL of 0.289 M HCl to titrate the excess NaOH. Calculate the purity of the aspirin. What indicator should be used for this titration? Why?

The pigment cyanidin aglycone is one of the anthocyanin molecules that gives red cabbage (Brassica oleracea var. capitata f. rubra) its characteristic red coloration. Many young chemists have used this red cabbage indicator to study acidbase chemistry. Estimate the pH range at which cyanidin aglycone shows a color change

Amino acids are the building blocks for all proteins in our bodies. A structure for the amino acid alanine is All amino acids have at least two functional groups with acidic or basic properties. In alanine, the carboxylic acid group has Ka 4.5  103 and the amino group has Kb 7.4  105 . Because of the two groups with acidic or basic properties, three different charged ions of alanine are possible when alanine is dissolved in water. Which of these ions would predominate in a solution with [H] 1.0 M? In a solution with [OH] 1.0 M?

Derive an equation analogous to the HendersonHasselbalch equation but relating pOH and pKb of a buffered solution composed of a weak base and its conjugate acid, such as NH3 and NH4 .

a. Calculate the pH of a buffered solution that is 0.100 M in C6H5CO2H (benzoic acid, Ka 6.4  105 ) and 0.100 M in C6H5CO2Na.

b. Calculate the pH after 20.0% (by moles) of the benzoic acid is converted to benzoate anion by addition of a strong base. Use the dissociation equilibrium to calculate the pH.

c. Do the same as in part b, but use the following equilibrium to calculate the pH:

d. Do your answers in parts b and c agree? Explain.

Consider a solution containing 0.10 M ethylamine (C2H5NH2), 0.20 M C2H5NH3 , and 0.20 M Cl. a. Calculate the pH of this solution. b. Calculate the pH after 0.050 mol KOH(s) is added to 1.00 L of this solution. (Ignore any volume changes.)

You make 1.00 L of a buffered solution (pH 4.00) by mixing acetic acid and sodium acetate. You have 1.00 M solutions of each component of the buffered solution. What volume of each solution do you mix to make such a buffered solution?

You have the following reagents on hand: C6H5CO2 1aq2 H2O1l2 C6H5CO2H1aq2 OH1aq2 C6H5CO2H1aq2 C6H5CO2 1aq2 H1aq2 Solids (pKa of Acid Form Is Given) Solutions Benzoic acid (4.19) 5.0 M HCl Sodium acetate (4.74) 1.0 M acetic acid (4.74) Potassium fluoride (3.14) 2.6 M NaOH Ammonium chloride (9.26) 1.0 M HOCl (7.46) What combinations of reagents would you use to prepare buffers at the following pH values? a. 3.0 b. 4.0 c. 5.0 d. 7.0 e. 9.0 88

What quantity (moles) of HCl(g) must be added to 1.0 L of 2.0 M NaOH to achieve a pH of 0.00? (Neglect any volume changes.)

Calculate the value of the equilibrium constant for each of the following reactions in aqueous solution.

The following plot shows the pH curves for the titrations of various acids by 0.10 M NaOH (all of the acids were 50.0-mL samples of 0.10 M concentration).

Calculate the volume of 1.50  102 M NaOH that must be added to 500.0 mL of 0.200 M HCl to give a solution that has pH 2.15.

Repeat the procedure in Exercise 55, but for the titration of 25.0 mL of 0.100 M HNO3 with 0.100 M NaOH.

A certain acetic acid solution has pH 2.68. Calculate the volume of 0.0975 M KOH required to reach the equivalence point in the titration of 25.0 mL of the acetic acid solution.

A 0.210-g sample of an acid (molar mass 192 g/mol) is titrated with 30.5 mL of 0.108 M NaOH to a phenolphthalein end point. Is the acid monoprotic, diprotic, or triprotic?

A student intends to titrate a solution of a weak monoprotic acid with a sodium hydroxide solution but reverses the two solutions and places the weak acid solution in the buret. After 23.75 mL of the weak acid solution has been added to 50.0 mL of the 0.100 M NaOH solution, the pH of the resulting solution is 10.50. Calculate the original concentration of the solution of weak acid

A student titrates an unknown weak acid, HA, to a pale pink phenolphthalein end point with 25.0 mL of 0.100 M NaOH. The student then adds 13.0 mL of 0.100 M HCl. The pH of the resulting solution is 4.7. How is the value of pKa for the unknown acid related to 4.7?

A sample of a certain monoprotic weak acid was dissolved in water and titrated with 0.125 M NaOH, requiring 16.00 mL to reach the equivalence point. During the titration, the pH after adding 2.00 mL NaOH was 6.912. Calculate Ka for the weak acid.

Another way to treat data from a pH titration is to graph the absolute value of the change in pH per change in milliliters added versus milliliters added ( pH mL versus mL added). Make this graph using your results from Exercise 55. What advantage might this method have over the traditional method for treating titration data?

A buffer is made using 45.0 mL of 0.750 M HC3H5O2 (Ka 1.3  105 ) and 55.0 mL of 0.700 M NaC3H5O2. What volume of 0.10 M NaOH must be added to change the pH of the original buffer solution by 2.5%?

A 0.400 M solution of ammonia was titrated with hydrochloric acid to the equivalence point, where the total volume was 1.50 times the original volume. At what pH does the equivalence point occur?

What volume of 0.0100 M NaOH must be added to 1.00 L of 0.0500 M HOCl to achieve a pH of 8.00?

Consider a solution formed by mixing 50.0 mL of 0.100 M H2SO4, 30.0 mL of 0.100 M HOCl, 25.0 mL of 0.200 M NaOH, 25.0 mL of 0.100 M Ba(OH)2, and 10.0 mL of 0.150 M KOH. Calculate the pH of this solution

When a diprotic acid, H2A, is titrated with NaOH, the protons on the diprotic acid are generally removed one at a time, resulting in a pH curve that has the following generic shape:

Consider the following two acids: HO2CCH2CH2CH2CH2CO2H Adipic acid In two separate experiments the pH was measured during the titration of 5.00 mmol of each acid with 0.200 M NaOH. Each experiment showed only one stoichiometric point when the data were plotted. In one experiment the stoichiometric point was at 25.00 mL added NaOH, and in the other experiment the stoichiometric point was at 50.00 mL NaOH. Explain these results (See Exercise 103.)

The titration of Na2CO3 with HCl has the following qualitative profile: a. Identify the major species in solution at points AF. b. Calculate the pH at the halfway points to equivalence, B and D. (Hint: Refer to Exercise 103.)

Consider the titration curve in Exercise 105 for the titration of Na2CO3 with HCl. a. If a mixture of NaHCO3 and Na2CO3 was titrated, what would be the relative sizes of V1 and V2? b. If a mixture of NaOH and Na2CO3 was titrated, what would be the relative sizes of V1 and V2? c. A sample contains a mixture of NaHCO3 and Na2CO3. When this sample was titrated with 0.100 M HCl, it took 18.9 mL to reach the first stoichiometric point and an additional 36.7 mL to reach the second stoichiometric point. What is the composition in mass percent of the sample?

A few drops of each of the indicators shown in the accompanying table were placed in separate portions of a 1.0 M solution of a weak acid, HX. The results are shown in the last column of the table. What is the approximate pH of the solution containing HX? Calculate the approximate value of Ka for HX.

Malonic acid (HO2CCH2CO2H) is a diprotic acid. In the titration of malonic acid with NaOH, stoichiometric points occur at pH 3.9 and 8.8. A 25.00-mL sample of malonic acid of unknown concentration is titrated with 0.0984 M NaOH, requiring 31.50 mL of the NaOH solution to reach the phenolphthalein end point. Calculate the concentration of malonic acid in the unknown solution. (See Exercise 103.)

A buffer solution is prepared by mixing 75.0 mL of 0.275 M fluorobenzoic acid (C7H5O2F) with 55.0 mL of 0.472 M sodium fluorobenzoate. The pKa of this weak acid is 2.90. What is the pH of the buffer solution?

A 10.00-g sample of the ionic compound NaA, where A is the anion of a weak acid, was dissolved in enough water to make 100.0 mL of solution and was then titrated with 0.100 M HCl. After 500.0 mL HCl was added, the pH was 5.00. The experimenter found that 1.00 L of 0.100 M HCl was required to reach the stoichiometric point of the titration. a. What is the molar mass of NaA? b. Calculate the pH of the solution at the stoichiometric point of the titration

Calculate the pH of a solution prepared by mixing 250. mL of 0.174 m aqueous HF (density 1.10 g/mL) with 38.7 g of an aqueous solution that is 1.50% NaOH by mass (density 1.02 g/mL). (Ka for HF 7.2  104 .)

A 225-mg sample of a diprotic acid is dissolved in enough water to make 250. mL of solution. The pH of this solution is 2.06. A 6.9  103 M solution of calcium hydroxide is prepared. Enough of the calcium hydroxide solution is added to the solution of the acid to reach the second equivalence point. The pH at the second equivalence point (as determined by a pH meter) is 7.96. The first dissociation constant for the acid is 5.90 1Ka1 2 102 . Assume that the volumes of the solutions are additive, that all solutions are at 25 C, and that is at least 1000 times greate than . a. Calculate the molar mass of the acid. b. Calculate the second dissociation constant for the acid .