Calculate the pH of a solution that is a. 0.0100 M in HClO4 and 0.0300 M in monochloroacetic acid. b. 0.0100 M in HCl and 0.0150 M in H2SO4. c. 0.0100 M in NaOH and 0.0300 M in Na2S. d. 0.0100 M in NaOH and 0.0300 M in sodium acetate.
> Suggest a range of sample masses for the indicated primary standard if it is desired to use between 35 and 45 mL of titrant a. 0.180 M HClO4 titrated against Na2CO3(CO2 product) b. 0.102 M HCl titrated against Na2C2O4 c. 0.180 M NaOH titrated against ben
> Calculate the molar concentration of a dilute BaOH2 solution if a. 50.00 mL yielded 0.1791 g of BaSO4 b. titration of 0.4512 g of primary-standard potassium hydrogen phthalate (KHP) required 26.46 mL of the base c. addition of 50.00 mL of the base to 0
> Calculate the molar concentration of a dilute HCl solution if a. A 50.00-ml aliquot yielded 0.6027 g of agcl. b. The titration of 25.00 ml of 004096mbaOH2 required 17.93 ml of the acid c. The titration of 0.2407 g of primary-standard Na2CO3 requir
> A NaOH solution was 0.1019 M immediately after standardization. Exactly 500.0 mL of the reagent was left exposed to air for several days and absorbed 0.652 g of CO2 Calculate the relative carbonate error in the determination of acetic acid with this so
> If 1.000 L of 0.2000 M NaOH was unprotected from the air after standardization and absorbed 14.2 mmol of CO2 , what is its new molar concentration when it is standardized against a standard solution of HCl using
> The concentration of a perchloric acid solution was established by titration against primary-standard sodium carbonate (product: CO2 ); the following data were obtained.
> Standardization of a sodium hydroxide solution against potassium hydrogen phthalate (KHP) yielded the results in the following table. a. the average molar concentration of the base. b. the standard deviation and the coefficient of variation for the dat
> A series of sulfate samples is to be analyzed by precipitation as BaSO4. If it is known that the sulfate content in these samples ranges between 20% and 55%, what minimum sample mass should be taken to ensure that a precipitate mass no smaller than 0.200
> How would you prepare 2.00 L of a. 0.10 M KOH from the solid? b. 0010MBaOH2 8H2O from the solid? c. 0.150 M HCl from a reagent that has a density of 1.0579 g/mL and is 11.50% HCl (w/w)?
> The mishandling of a shipping container loaded with 850 cases of wine caused some of the bottles to break. An insurance adjuster proposed to settle the claim at 20.8% of the value of the shipment, based on a random 250-bottle sample in which 52 were crac
> How would you prepare 500 mL of a. 0100MHSO24 from a reagent that has a density of 1.1539 g/mL and is 21.8% H2SO4w/w? b. 0.200 M NaOH from the solid? c. 008000MNa2CO3 from the pure solid?
> The solubility products for a series of hydroxides are Which hydroxide has a. the lowest molar solubility in H2O? b. the lowest molar solubility in a solution that is 0.35 M in NaOH?
> Briefly describe the circumstance where the concentration of a sodium hydroxide solution will apparently be unaffected by the absorption of carbon dioxide.
> What types of organic nitrogen-containing compounds tend to yield low results with the Kjeldahl method unless special precautions are taken?
> Suggest an indicator that could be used to provide an end point for the titration of the first proton in H3AsO4.
> Indicate whether an aqueous solution of the following compounds is acidic, neutral, or basic. Explain your answer. a. NaH2PO4 b. Na3PO4 c. NaNO3 d. NaHC2O4 e. Na2C2O4 f. Na2HPO4 g. NH4OAc h. NaNO2
> Why is it impossible to titrate all three protons of phosphoric acid in aqueous solution?
> Briefly explain why Equation 13-15 can only be used to calculate the hydronium ion concentration of solutions in which NaHA is the only solute that determines the pH.
> Find the number of millimoles of solute in a. 2.00 L of 0.0449 M KMnO4. b. 750 mL of 5.35 × 10-3 M KSCN. c. 3.50 L of a solution that contains 6.23 ppm of CuSO4. d. 250 mL of 0.414 mM KCl.
> Explain the origin and significance of each of the terms on the right side of Equation 13-12. Does the equation make intuitive sense? Why or why not?
> As its name implies, NaHA is an “acid salt” because it has a proton available to donate to a base. Briefly explain why a pH calculation for a solution of NaHA differs from that for a weak acid of the type HA.
> Find the number of Na+ ions in 2.75 g of Na3PO4?
> Calculate alpha values for the following triprotic acids every 0.5 pH units from pH 0.0 to 14.0. Plot the distribution diagram for each of the acids, and label the curve for each species. a. Citric acid b. Arsenic acid
> Calculate alpha values for the following diprotic acids every 0.5 pH units from pH 0.0 to 10.0. Plot the distribution diagram for each of the acids, and label the curve for each species. a. Phthalic acid b. Succinic acid c. Tartaric acid
> Calculate the p-value for each of the listed ions in the following: a. Na1, Cl2, and OH2 in a solution that is 0.0635 M in NaCl and 0.0403 M in NaOH. b. Ba2+, Mn2+, and Cl2 in a solution that is 4.65 ×10-3 M in BaCl2 and 2.54 M in MnCl2. c. H+, Cl-, and
> Derive equations that define a0, a1, a2, and a3 for the acid H3AsO4.
> For pH values of 2.00, 6.00, and 10.00, calculate the alpha values for each species in an aqueous solution of a. Phthalic acid. b. Tartaric acid. c. Citric acid. d. Periodic acid. e. Phosphorous acid. f. Malonic acid.
> Calculate a numerical value for the equilibrium constant for the reaction
> Formulate equilibrium constants for the following equilibria, and determine numerical values for the constants: a. b.
> A 0.6447-g portion of manganese dioxide was added to an acidic solution in which 1.1402 g of a chloride containing sample was dissolved. Evolution of chlorine took place as a consequence of the following reaction: After the reaction was complete, the ex
> Generate a curve for the titration of 50.00 mL of a solution in which the analytical concentration of HClO4 is 0.1000 M and that for formic acid is 0.0800 M. Calculate the pH after addition of 0.00, 10.00, 20.00, 24.00, 25.00, 26.00, 35.00, 44.00, 45.00,
> Generate a curve for the titration of 50.00 mL of a solution in which the analytical concentration of NaOH is 0.1000 M and that for hydrazine is 0.0800 M. Calculate the pH after addition of 0.00, 10.00, 20.00, 24.00, 25.00, 26.00, 35.00, 44.00, 45.00, 46
> Construct a curve for the titration of 50.00 mL of a 0.1000 M solution of compound A with a 0.2000 M solution of compound B in the following table. For each titration, calculate the pH after the addition of 0.00, 12.50, 20.00, 24.00, 25.00, 26.00, 37.50,
> Changes in the method used to coat the tablets in Problem 6-6 lowered the percentage of rejects from 8.0% to 3.0%. How many tablets should be taken for inspection if the permissible relative standard deviation in the measurement is to be a. 15%?
> Briefly explain why curve B cannot describe the titration of a mixture consisting of H3PO4 and NaH2PO4.
> Describe the composition of a solution that would be expected to yield a curve resembling (see Problem 13-26) a. curve B. b. curve A. c. curve E.
> Identify by letter the curve you would expect in the titration of a solution containing a. Disodium maleate, Na2M, with standard acid. b. Pyruvic acid, HP, with standard base. c. Sodium carbonate, Na2CO3, with standard acid.
> The solubility products for a series of iodides are List these four compounds in order of decreasing molar solubility in (a) water. (b) 0.20 M NaI. (c) a 0.020 M solution of the solute cation.
> How would you prepare 1.00 L of a buffer with a pH of 6.00 from 0.500 M Na3AsO4 and 0.400 M HCl?
> How would you prepare 1.00 L of a buffer with a pH of 7.00 from 0.200 M H3PO4 and 0.160 M NaOH?
> Ammoniacal nitrogen can be determined by treatment of the sample with chloroplatinic acid; the product is slightly soluble ammonium chloroplatinate: The precipitate decomposes on ignition, yielding metallic platinum and gaseous products: Calculate the pe
> How would you prepare 1.00 L of a buffer with a pH of 9.45 from 0.300 M Na2CO3 and 0.200 M HCl?
> What is the pH of the buffer formed by adding 100 mL of 0.150 M potassium hydrogen phthalate to a. 100.0 mL of 0.0800 M NaOH? b. 100.0 mL of 0.0800 M HCl?
> . What is the pH of the buffer formed by mixing 50.0 mL of 0.200 M NaH2PO4 with a. 50.0 mL of 0.100 M HCl? b. 50.0 mL of 0.100 M NaOH?
> What mass (g) of dipotassium phthalate must be added to 750 mL of 0.0500 M phthalic acid to give a buffer of pH 5.75?
> A coating that weighs at least 2.00 mg is needed to impart adequate shelf life to a pharmaceutical tablet. A random sampling of 200 tablets revealed that 16 failed to meet this requirement. a. Use this information to estimate the relative standard deviat
> What mass (g) of Na2HPO4 . 2H2O must be added to 750 mL of 0.160 M H3PO4 to give a buffer of pH 7.30?
> Identify the principal conjugate acid-base pair and calculate the ratio between them in a solution that is buffered to pH 9.00 and contains a. H2S. b. ethylenediamine dihydrochloride. c. H3AsO4. d. H2CO3.
> Identify the principal conjugate acid-base pair and calculate the ratio between them in a solution that is buffered to pH 6.00 and contains a. H2SO3. b. Citric acid. c. Malonic acid. d. Tartaric acid.
> What is the mass in grams of solute in a. 250 mL of 0.264 M H2O2? b. 37.0 mL of 5.75 × 10-4 M benzoic acid (122 g/mol)? c. 2.50 L of a solution that contains 37.2 ppm of SnCl2? d. 11.7 mL of 0.0225 M KBrO3?
> The iodide in a sample that also contained chloride was converted to iodate by treatment with an excess of bromine: The unused bromine was removed by boiling; an excess of barium ion was then added to precipitate the iodate: In the analysis of a 1.59-g
> Calculate the pH of a solution that is a. 0.0100 M in HCl and 0.0200 M in picric acid. b. 0.0100 M in HCl and 0.0320 M in benzoic acid. c. 0.0100 M in NaOH and 0.075 M in Na2CO3. d. 0.0100 M in NaOH and 0.090 M in NH3.
> Calculate the pH of a solution that contains the following analytical concentrations: a. 0.225 M in H3PO4 and 0.414 M in NaH2PO4. b. 0.0670 M in Na2SO3 and 0.0315 M in NaHSO3. c. 0.640 M in HOC2H4NH2 and 0.750 M in HOC2H4NH3Cl. d. 0.0240 in H2C2O4 (o
> Calculate the pH of a solution that contains the following analytical concentrations: a. 0.0200 M in H3PO4 and 0.0500 M in NaH2PO4. b. 0.0300 M in NaH2AsO4 and 0.0500 M in Na2HAsO4. c. 0.0400 M in Na2CO3 and 0.0500 M in NaHCO3. d. 0.0400 M in H3P
> Calculate the pH of a solution that is 0.0400 M in a. Na3PO4. b. Na2C2O4. c. Na2HPO3. d. Na2SO3. e. Na2S. f. C2H4(NH3 1CI2)2.
> Calculate the pH of a solution that is 0.0400 M in a. NaH2PO4. b. NaHC2O4. c. NaH2PO3. d. NaHSO3. e. NaHS. f. H2NC2H4NH3 +CI-.
> The following results were obtained for the determination of calcium in a NIST limestone sample: %CaO = 51.33, 51.22, 51.36, 51.21, and 51.44. Five gross samples were then obtained for a carload of limestone. The average percent CaO values for the gross
> Calculate the pH of a solution that is 0.0400 M in a. H3PO4. b. H2C2O4. c. H3PO3. d. H2SO3. e. H2S. f. H2NC2H4NH2.
> Suggest a suitable indicator for a titration based on each of the following reactions. Use 0.05 M if an equivalence point concentration is needed. a. H2CO3 + NaOH→ NaHCO3 + H2O b. H2P + 2NaOH → Na2P + 2H2O (H2P 5 o@phthalic acid) c. H2T + 2NaOH → Na2
> Suggest a method for determining the amounts of H3PO4 and NaH2PO4 in an aqueous solution.
> Suggest an indicator that would give an end point for the titration of the first two protons in H3AsO4.
> The mercury in a 1.0451-g sample was precipitated with an excess of paraperiodic acid, H5IO6: The precipitate was filtered, washed free of precipitating agent, dried, and weighed, and 0.5718 g was recovered. Calculate the percentage of Hg2Cl2 in the sa
> Why are the standard reagents used in neutralization titrations generally strong acids and bases rather than weak acids and bases?
> The solubility-product constant for K2PdCl6 is What is the K1 concentration of a solution prepared by mixing 50.0 mL of 0.200 M KCl with 50.0 mL of a. 0.0800 M PdCl62-? b. 0.160 M PdCl62-? c. 0.240 M PdCl62-?
> Consider curves for the titration of 0.10 M NaOH and 0.010 M NH3 with 0.10 M HCl. a. Briefly account for the differences between curves for the two titrations. b. In what respect will the two curves be indistinguishable?
> What factors affect end-point sharpness in an acid-base titration?
> Why does the typical acid-base indicator exhibit its color change over a range of about 2 pH units?
> Supply the missing data in the following table.
> What factors determine the mass of a gross sample?
> Calculate the molar concentration of a solution that is 50.0% NaOH (w/w) and has a specific gravity of 1.52..
> Calculate the equilibrium concentration of methyl ammonia in a solution that has a molar analytical CH3NH2 concentration of 0.120 and a pH of 11.471.
> Calculate a0 and a1 for a. Acetic acid in a solution with a ph of 5.320. b. Picric acid in a solution with a ph of 1.750. c. Hypochlorous acid in a solution with a ph of 7.00. d. Hydroxylamine acid in a solution with a ph of 5.54. e. Piperidine in a solu
> Five different laboratories participated in an interlaboratory study involving determinations of the iron level in water samples. The following results are replicate determinations of Fe in ppm for laboratories A through E. a. State the appropriate hypot
> Calculate the pH after addition of 0.00, 5.00, 15.00, 25.00, 40.00, 49.00, 50.00, 51.00, 55.00, and 60.00 mL of reagent in the titration of 50.0 mL of a. 0.01000 M chloroacetic acid with 0.01000 M NaOH. b. 0.1000 M anilinium chloride with 0.1000 M NaO
> Calculate the pH after addition of 0.00, 5.00, 15.00, 25.00, 40.00, 45.00, 49.00, 50.00, 51.00, 55.00, and 60.00 mL of 0.1000 M HCl in the titration of 50.00 mL of a. 0.1500 M ammonia. b. 0.1500 M hydrazine. c. 0.1500 M sodium cyanide.
> Calculate the pH after addition of 0.00, 5.00, 15.00, 25.00, 40.00, 45.00, 49.00, 50.00, 51.00, 55.00, and 60.00 mL of 0.1000 M NaOH in the titration of 50.00 mL of a. 0.1000 M HNO2. b. 0.1000 M pyridinium chloride. c. 0.1000 M lactic acid.
> What is the mass in milligrams of solute in a. 16.0 mL of 0.350 M sucrose (342 g/mol)? b. 1.92 L of 3.76 × 10-3 M H2O2? c. 356 mL of a solution that contains 2.96 ppm of Pb (NO3)2? d. 5.75 mL of 0.0819 M KNO3?
> Briefly describe or define and give an example of a. An amphiprotic solvent. b. A differentiating solvent. c. A leveling solvent. d. A mass-action effect.
> In a titration of 50.00 mL of 0.1000 M ethylamine with 0.1000 M HClO4, the titration error must be no more than 0.05 mL. What indicator can be chosen to realize this goal?
> In a titration of 50.00 mL of 0.05000 M formic acid with 0.1000 M KOH, the titration error must be smaller than 0.05 mL. What indicator can be chosen to realize this goal?
> Assuming a large number of measurements so that s is a good estimate of σ, determine what confidence level was used for each of the following confidence intervals.
> Before agreeing to the purchase of a large order of solvent, a company wants to see conclusive evidence that the mean value of a particular impurity is less than 1.0 ppb. What hypotheses should be tested? What are the type I and type II errors in this si
> A 50.00-mL aliquot of 0.1000 M NaOH is titrated with 0.1000 M HCl. Calculate the pH of the solution after the addition of 0.00, 10.00, 25.00, 40.00, 45.00, 49.00, 50.00, 51.00, 55.00, and 60.00 mL of acid, and prepare a titration curve from the data.
> A 5.500-g sample of a pesticide was decomposed with metallic sodium in alcohol, and the liberated chloride ion was precipitated as AgCl. Express the results of this analysis in terms of percent DDT 1C14H9Cl52 based on the recovery of 0.1873 g of AgCl.
> Calculate the change in pH to three decimal places that occurs when 0.50 mmol of a strong acid is added to 100 mL of a. 0.0100 M lactic acid 1 0.0800 M sodium lactate. b. 0.0800 M lactic acid 1 0.0200 M sodium lactate. c. 0.0500 M lactic acid 1 0.0500
> Calculate the change in pH that occurs when 1.50 mmol of a strong base is added to 100 mL of the solutions listed in Problem 12-34. Calculate values to three decimal places.
> Calculate the change in pH that occurs when 1.00 mmol of a strong acid is added to 100 mL of the solutions listed in Problem 12-3
> Calculate the change in pH that occurs in each of the following solutions as a result of a tenfold dilution with water. Round calculated values for pH to three figures to the right of the decimal point. a. H2O b. 0.0500 M HCl c. 0.0500 M NaOH d. 0.
> What is the pH of a solution that is a. Prepared by dissolving 3.30 g of 1NH422SO4 in water, adding 125.0 ml of 0.1011 M naoh, and diluting to 500.0 ml? b. 0.120 M in piperidine and 0.010 M in its chloride salt? c. 0.050 M in ethylamine and 0.167 M in it
> What is the pH of a solution that is a. Prepared by dissolving 6.75 g of lactic acid (90.08 g/mol) and 5.19 g of sodium lactate (112.06 g/mol) in water and diluting to 1.00 L? b. 0.0430 M in acetic acid and 0.0175 M in sodium acetate? c. Prepared by diss
> A solution is 0.0500 M in NH4Cl and 0.0300 M in NH3. Calculate its OH-concentration and its pH a.Neglecting activities. b.Taking activities into account.
> The solubility-product constant for Ce(IO3)3 is 3.2 × 10-10. What is the Ce3+ concentration in a solution prepared by mixing 50.00 mL of 0.0500 M Ce3+with 50.00 mL of a. Water? b. 0.0500 M IO3-? c. 0.250 M IO3-? d. 0.0450 M IO3-?
> Calculate the pH of the solution that results when 20.0 mL of 0.2500 M NH3 is a.Mixed with 20.0 ml of distilled water. b.Mixed with 20.0 ml of 0.250 M hcl solution. c.Mixed with 20.0 ml of 0.300 M hcl solution. d.Mixed with 20.0 ml of 0.200 M NH4Cl solut
> Describe in your own words why the confidence interval for the mean of five measurements is smaller than that for a single result.
> A 0.2121-g sample of an organic compound was burned in a stream of oxygen, and the CO2 produced was collected in a solution of barium hydroxide. Calculate the percentage of carbon in the sample if 0.6006 g of BaCO3 was formed.
> The ascorbic acid concentration in mmol/L of five different brands of orange juice was measured. Six replicate samples of each brand were analyzed. The following partial ANOVA table was obtained. a. Fill in the missing entries in the table. b. State the
> Sewage and industrial pollutants dumped into a body of water can reduce the dissolved oxygen concentration and adversely affect aquatic species. In one study, weekly readings are taken from the same location in a river over a two-month period. Some scien
> A prosecuting attorney in a criminal case presented as principal evidence small fragments of glass found embedded in the coat of the accused. The attorney claimed that the fragments were identical in composition to a rare Belgian stained glass window bro
> To test the quality of the work of a commercial laboratory, duplicate analyses of a purified benzoic acid (68.8% C, 4.953% H) sample were requested. It is assumed that the relative standard /reported results are 68.5% C and 4.882% H. At the 95% confidenc
> A standard method for the determination of glucose in serum is reported to have a standard deviation of 0.36 mg/dL. If s = 0.36 is a good estimate of σ, how many replicate determinations should be made in order for the mean for the analysis of a sample t
