Write balanced equations showing how 2 2 K Cr O could be used as a primary standard for solutions of 7Na S O 2+3.
> Find the number of millimoles of solute in a. 386 mL of 0.210 M HClO4. b. 25.0 L of 8.05 × 10-3 M K2CrO4. c. 4.50 L of an aqueous solution that contains 6.95 ppm of AgNO3. d. 537 mL of 0.0200 M KOH.
> What mass of solid PbCl2 (278.10 g/mol) is formed when 200 mL of 0.125 M Pb2+ is mixed with 400 mL of 0.175 M Cl?
> What mass of solid La(IO3)3 (663.6 g/mol) is formed when 50.0 mL of 0.250 M La3+ is mixed with 75.0 mL of 0.302 M IO3-?
> Describe the preparation of a. 2.50 L of 0.0250 M KMnO4 from the solid reagent. b. 4.00 L of 0.250 M HClO4, starting with an 8.00 M solution of the reagent. c. 500 mL of a solution that is 0.0200 M in I-, starting with MgI2. d. 200 mL of 1.00% (w/v) aque
> Describe the preparation of a. 500 mL of 0.1000 M AgNO3 from the solid reagent. b. 1.00 L of 0.1000 M HCl, starting with a 6.00 M solution of the reagent. c. 250 mL of a solution that is 0.0810 M in K+, starting with solid K4Fe(CN)6. d. 500 mL of 3.00% (
> Is the following statement true or false or both? Define your answer with equations, examples, or graphs. “A buffer maintains the pH of a solution constant.”
> Describe the preparation of 750 mL of 3.00 M HNO3 from the commercial reagent that is 70.5% HNO3 (w/w) and has a specific gravity of 1.42.
> Give two examples of units derived from the fundamental base SI units.
> What volume of 2.00 M NaOH must be added to 200.0 mL of 1.00 M glycolic acid to produce a buffer solution having a pH of 4.15?
> Describe the preparation of 500 mL of 3.00 M H3PO4from the commercial reagent that is 86% H3PO4(w/w) and has a specific gravity of 1.71.
> The following table gives the sample means and standard deviations for six measurements each day of the purity of a polymer in a process. The purity is monitored for 24 days. Determine the overall mean and standard deviation of the measurements, and cons
> What mass of solute in milligrams is contained in a. 26.0 mL of 0.250 M sucrose (342 g/mol)? b. 2.92 L of 5.23 3 1024 M H2O2? c. 673 mL of a solution that contains 5.76 ppm Pb1NO322 1331.20 g>mol2? d. 6.75 mL of 0.0426 M KNO3?
> What mass of KIO3 is needed to convert the copper in 0.1570 g of CuSO4 . 5H2O to Cu1IO322?
> What volume of 0.200 M HCl must be added to 500.0 mL of 0.300 M sodium mandelate to produce a buffer solution with a pH of 3.25?
> Describe the preparation of a. 1.50 L of 21.0% (w/v) aqueous glycerol (C3H8O3, 92.1 g/mol). b. 1.50 kg of 21.0% (w/w) aqueous glycerol. c. 1.50 L of 21.0% (v/v) aqueous glycerol.
> What mass of sodium glycolate should be added to 400.0 mL of 1.00 M glycolic acid to produce a buffer solution with a pH of 4.25?
> Describe the preparation of a. 500 mL of 5.25% (w/v) aqueous ethanol (C2H5OH, 46.1 g/mol). b. 500 g of 5.25% (w/w) aqueous ethanol. c. 500 mL of 5.25% (v/v) aqueous ethanol.
> What mass of sodium formate must be added to 500.0 mL of 1.00 M formic acid to produce a buffer solution that has a pH of 3.75?
> Why are Ce4+ solutions never used for the titration of reductants in basic solutions?
> Why are Ce4+ solutions never used for the titration of reductants in basic solutions?
> Why are standard solutions of reductants less often used for titrations than standard solutions of oxidants?
> Write a balanced net ionic equation for the reduction of UO22+ in a Walden reductor.
> Why is a Walden reductor always used with solutions that contain appreciable concentrations of HCl?
> How many millimoles of solute are contained in a. 2.95 mL of 0.0789 M KH2PO4? b. 0.2011 L of 0.0564 M HgCl2? c. 2.56 L of a 47.5 ppm solution of Mg1NO322? d. 79.8 mL of 0.1379 M NH4VO3 1116.98 g>mol2?
> The following data represent measurements made on a process for 30 days. One measurement was made each day. Assuming that 30 measurements are enough that x S m and s S s, find the mean of the values, the standard deviation, and the upper and lower contro
> Write balanced net ionic equations to describe
> Use a spreadsheet to do the calculations and plot the titration curves for the following titrations. Calculate potentials after the addition of titrant corresponding to 10%, 20%, 30%, 40%, 50%, 60%, 70%, 80%, 90%, 95%, 99%, 99.9%, 100%, 101%, 105%, 110%,
> A 11.4% (w/w) NiCl2 (129.61 g/mol) solution has a density of 1.149 g/mL. Calculate a. The molar concentration of NiCl2 in this solution. b. The molar Cl2 concentration of the solution. c. The mass in grams of NiCl2 contained in each liter of this solutio
> A gas mixture was passed at the rate of 2.50 L/min through a solution of sodium hydroxide for a total of 59.00 min. The 2 SO in the mixture was retained as sulfite ion: After acidification with HCl, the sulfite was titrated with 5.15 mL of 0.002997MKIO3
> A 2.552-g sample containing both Fe and V was dissolved under conditions that converted the elements to Fe(III) and V(V). The solution was diluted to 500.0 mL, and a 50.00-mL aliquot was passed through a Walden reductor and titrated with 17.79 mL of 4 0
> A sensitive method for I− in the presence of CI− and Br− entails oxidation of the I− to IO3−with Br2. The excess Br2 is then removed by boiling or by reduction with formate ion. The IO3− produced is determined by addition of excess I- and titration of th
> The ethyl mercaptan concentration in a mixture was determined by shaking a 1.795-g sample with 50.00 mL of 2 0 01204 MI. in a tightly stoppered flask: The excess 2I was back-titrated with 15.21 mL of 22 3 0 01437MNa S O . . Calculate thepercentage of 2 5
> An 8.13-g sample of an ant-control preparation was decomposed by wet ashing with H2SO4 and HNO3. The As in the residue was reduced to the trivalent state with hydrazine. After removal of the excess reducing agent, the As(III) required a 31.46-mL titratio
> The 3 KClO in a 0.1791-g sample of an explosive was determined by reaction with 50.00 mL of 0 0. 873 2 MFe + When the reaction was complete, the excess 2 Fe + was back-titrated with 14.95 mL of 4 0 06970MCe + . . Calculate the percentage of 3 KClO in the
> How many millimoles of solute are contained in a. 2.00 L of 2.76 3 1023 M KMnO4? b. 250.0 mL of 0.0423 M KSCN? c. 500.0 mL of a solution containing 2.97 ppm CuSO4? d. 2.50 L of 0.352 M KCl?
> Treatment of hydroxylamine 2 (H NOH) with an excess of Fe(III) results in the formation of N O2 and an equivalent amount of Fe(II): Calculate the molar concentration of an H N2 OH solution if the Fe(II) produced by treatment of a 25.00-mL aliquot requir
> Atomic emission measurements were made to determine sodium in a blood serum sample. The following emission intensities were obtained for standards of 5.0 and 10.0 ng/mL and for the serum sample. All emission intensities were corrected for any blank emiss
> There is 0.5690-g specimen of iron ore was dissolved and passed through a Jones reductor. Titration of the Fe(II) produced required 38.79 mL of 0.01926 M KMnO4. Express the results of this analysis in terms of (a) percent Fe and (b) percent 2 3 Fe O .
> Calculate the percentage of MnO2 in a mineral specimen if the 2I liberated by a 0.1267-g sample in the net reaction
> A 0.1853-g sample of 3 KBrO was dissolved in dilute HCl and treated with an unmeasured excess of KI. The liberated iodine required 44.36 mL of a sodium thiosulfate solution. Calculate the molar concentration of the Na S O 22 3.
> Consult Appendix 3, and pick out a suitable acid-base pair to prepare a buffer with a pH of a. 10.3. b. 6.1. c. 4.5. d. 8.1.
> A 0.2219-g sample of pure iron wire was dissolved in acid, reduced to the +2 state, and titrated with 34.65 mL of cerium (IV). Calculate the molar concentration of the 4 Ce +solution.
> How would you prepare 1.000 L of 0.05000 M KBrO3?
> How would you prepare 2.0 L of approximately 0.04 M I3 solution? Calculate the molar concentration of 4 KMnO in this solution.
> Why are KMnO4 solutions filtered before they are standardized?
> Calculations of volumetric analysis ordinarily consist of transforming the quantity of titrant used (in chemical units) to a chemically equivalent quantity of analyte (also in chemical units) through use of a stoichiometric factor. Use chemical formulas
> A solution prepared by dissolving a 0.2541-g sample of electrolytic iron wire in acid was passed through a Jones reductor. The iron(II) in the resulting solution required a 36.76-mL titration. Calculate the molar oxidant concentration if the titrant used
> In the titration of 2I solutions with Na S O 2+ 3 , the starch indicator is never added until just before chemical equivalence. Why?
> The following atomic absorption results were obtained for determinations of Zn in multivitamin tablets. All absorbance values are corrected for the appropriate reagent blank (cZn = 0.0 ng/mL). The mean value for the blank was 0.0000 with a standard devia
> Suggest a way in which a solution of 3 KIO could be used as a source of known quantities of 2I .
> A standard solution of 2I increased in concentration with standing. Write a balanced net ionic equation that accounts for the increase.
> What is the primary use of standard K2Cr2O7 solutions?
> A 5.85% (w/w) Fe(NO3)3 (241.86 g/mol) solution has a density of 1.059 g/mL. Calculate a. The molar analytical concentration of Fe(NO3)3 in this solution. b. The molar NO3- concentration in the solution. c. The mass in grams of Fe(NO3)3 contained in each
> Briefly explain why there is no term in an equilibrium constant expression for water or for a pure solid, even though one (or both) appears in the balanced net ionic equation for the equilibrium.
> Why are solutions of KMnO4 and Na2S2O3 generally stored in dark reagent bottles? Answer Standard permanganate and thiosulfate solutions are generally stored in the dark because their decomposition reactions are catalyzed by light.
> Briefly explain why the concentration units of milligrams of solute per liter and parts per million can be used interchangeably for a dilute aqueous solution.
> Why are KMnO4 solutions filtered before they are standardized?
> Under what circumstance is the curve for an oxidation/reduction titration asymmetric about the equivalence point?
> How does calculation of the electrode potential of the system at the equivalence point differ from that for any other point of an oxidation/reduction titration?
> The method of standard additions was used to determine nitrite in a soil sample. A 1.00-mL portion of the sample was mixed with 24.00 mL of a colorimetric reagent, and the nitrite was converted to a colored product that produced a blank-corrected absorba
> How is an oxidation/reduction titration curve generated through the use of standard electrode potentials for the analyte species and the volumetric titrant?
> For an oxidation/reduction titration, briefly distinguish between a. Equilibrium and equivalence. b. A true oxidation/reduction indicator and a specific indicator.
> Use a spreadsheet, and construct curves for the following titrations. Calculate potentials after the addition of 10.00, 25.00, 49.00, 49.90, 50.00, 50.10, 51.00, and 60.00 mL of the reagent. Where necessary, assume that [H+] = 1.00 throughout.
> Select an indicator from Table 17-3 that might be suitable for each of the titrations in Problem 17-11. Write NONE if no indicator listed in Table 17-3 is suitable.
> An aqueous solution contains NaNO3 and KBr. The bromide ion is precipitated as AgBr by addition of AgNO3. After an excess of the precipitating reagent has been added, a. What is the charge on the surface of the coagulated colloidal particles? b. What i
> Consider solutions prepared by a. Dissolving 8.00 mmol of NaOAc in 200 mL of 0.100 M HOAc. b. Adding 100 mL of 0.0500 M NaOH to 100 mL of 0.175 M HOAc. c. Adding 40.0 mL of 0.1200 M HCl to 160.0 mL of 0.0420 M NaOAc. In what respects do these solution
> If you start with 0.1000 M solutions and the first-named species is the titrant, what will be the concentration of each reactant and product at the equivalence point of titrations (a), (c), (f), and (g) in Problem 17-11? Assume that there is no change in
> If you start with 0.1000 M solutions and the first-named species is the titrant, what will be the concentration of each reactant and product at the equivalence point of titrations (a), (c), (f), and (g) in Problem 17-11? Assume that there is no change in
> Calculate the electrode potential of the system at the equivalence point for each of the reactions in Problem 17-11. Use 0.100 M where a value for [H+] is needed and is not otherwise specified.
> Copper was determined in a river water sample by atomic absorption spectrometry and the method of standard additions. For the addition, 100.0μL of a 1000.0-μg/mL Cu standard was added to 100.0 mL of solution. The following data were obtained: Absorbance
> Generate equilibrium-constant expressions for the following reactions. Calculate numerical values for Keq.
> Use the shorthand notation (page 415) to describe the cells in Problem 17-9. Each cell is supplied with a salt bridge to provide electrical contact between the solutions in the two cell compartments.
> Calculate the potential of the following two half-cells that are connected by a salt bridge: a. A galvanic cell consisting of a lead electrode (right electrode) immersed in 0.0250 M Pb2+ and a zinc electrode in contact with 0.1000 M Zn2+. b. A galvanic c
> Calculate the theoretical cell potential of the following cells. If the cell is short-circuited, indicate the direction of the spontaneous cell reaction. Zn | Zn2+ (0.1000 M) || Co2+ (5.87 × 10−4 M) | Co Pt| Fe3+ (0.1600
> Calculate the theoretical potential of the following cells. Indicate whether the reaction will proceed spontaneously in the direction considered (oxidation on the left; reduction on the right) or whether an external voltage source is needed to force this
> Four different fluorescence flow cell designs were compared to see if they were significantly different. The following results represented relative fluorescence intensities for four replicate measurements: a. State the appropriate hypotheses. b. Do the
> Under what circumstance is the curve for an oxidation/reduction titration asymmetric about the equivalence point?
> How does calculation of the electrode potential of the system at the equivalence point differ from that for any other point of an oxidation/reduction titration?
> A solution was prepared by dissolving 875 mg of K3Fe(CN)6 (329.2 g/mol) in sufficient water to give750 mL. Calculate a. The molar analytical concentration of K3Fe(CN)6. b. The molar concentration of K+. c. The molar concentration of Fe(CN)3-. d. The weig
> How is an oxidation/reduction titration curve generated through the use of standard electrode potentials for the analyte species and the volumetric titrant?
> What is unique about the condition of equilibrium in an oxidation/reduction reaction?
> Potassium can be determined by flame emission spectrometry (flame photometry) using a lithium internal standard. The following data were obtained for standard solutions of KCl and an unknown containing a constant, known amount of LiCl as the internal sta
> The level of a pollutant in a river adjacent to a chemical plant is regularly monitored. Over a period of years, the normal level of the pollutant has been established by chemical analyses. Recently, the company has made several changes to the plant that
> For an oxidation/reduction titration, briefly distinguish between a. equilibrium and equivalence. b. a true oxidation/reduction indicator and a specific indicator.?
> Briefly define the electrode potential of a system that contains two or more redox couples.
> Use a spreadsheet, and construct curves for the following titrations. Calculate potentials after the addition of 10.00, 25.00, 49.00, 49.90, 50.00, 50.10, 51.00, and 60.00 mL of the reagent. Where necessary, assume that [H+] = 1.00 throughout.
> How can the relative supersaturation be varied during precipitate formation?
> Use a spreadsheet, and construct curves for the following titrations. Calculate potentials after the addition of 10.00, 25.00, 49.00, 49.90, 50.00, 50.10, 51.00, and 60.00 mL of the reagent. Where necessary, assume that [H1] 5 1.00 throughout. a. 50.00
> Select an indicator from Table 17-3 that might be suitable for each of the titrations in Problem 17-11. Write NONE if no indicator listed in Table 17-3 is suitable.
> If you start with 0.1000 M solutions and the first-named species is the titrant, what will be the concentration of each reactant and product at the equivalence point of titrations (a), (c), (f), and (g) in Problem 17-11? Assume that there is no change in
> Calculate the electrode potential of the system at the equivalence point for each of the reactions in Problem 17-11. Use 0.100 M where a value for [H+] is needed and is not otherwise specified.
> Generate equilibrium-constant expressions for the following reactions. Calculate numerical values for Keq.
> Which has the greater buffer capacity: (a) a mixture containing 0.100 mol of NH3 and 0.200 mol of NH4Cl or (b) a mixture containing 0.0500 mol of NH3 and 0.100 mol of NH4Cl?
> Use the shorthand notation (page 415) to describe the cells in Problem 17-9. Each cell is supplied with a salt bridge to provide electrical contact between the solutions in the two cell compartments.
> Water can be determined in solid samples by infrared spectroscopy. The water content of calcium sulfate hydrates is to be measured using calcium carbonate as an internal standard to compensate for some systematic errors in the procedure. A series of stan
> Calculate the potential of the following two half-cells that are connected by a salt bridge: a. a galvanic cell consisting of a lead electrode (right electrode) immersed in 0.0250 M Pb21 and a zinc electrode in contact with 0.1000 M Zn21. b. a galvanic
> Calculate the theoretical cell potential of the following cells. If the cell is short-circuited, indicate the direction of the spontaneous cell reaction
> Define what constitutes a chelating agent.
> Calculate the theoretical potential of the following cells. Indicate whether the reaction will proceed spontaneously in the direction considered (oxidation on the left; reduction on the right) or whether an external voltage source is needed to force this
> Why is it necessary to bubble hydrogen through the electrolyte in a hydrogen electrode?
