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.
> 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
> 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.
> 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?
> 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?
> The following entries are found in a table of standard electrode potentials: What is the significance of the difference between these two standard potentials?
> Make a clear distinction between a. oxidation and oxidizing agent. b. an electrolytic cell and a galvanic cell. c. the cathode of an electrochemical cell and the right-hand electrode. d. a reversible electrochemical cell and an irreversible electro
> Briefly describe or define a. electrode potential. b. formal potential. d. standard electrode potential. d. liquid junction potential. e. oxidation potential.
> Briefly describe or define a. oxidation. b. liquid junction. c. salt bridge. d. reductant. e. Nernst equation.
> Plot the half-cell potential versus concentration ratio for the half-cells of Problems 16-27 and 16-28. How would the plot look if potential were plotted against log(concentration ratio)?
> A solution was prepared by dissolving 5.76 g of KCl .MgCl2 . 6H2O (277.85 g/mol) in sufficient water to give 2.000 L. Calculate a. The molar analytical concentration of KCl . MgCl2 in this solution. b. The molar concentration of Mg2+. c. The molar concen
> A study was made to determine the activation energy EA for a chemical reaction. The rate constant k was determined as a function of temperature T, and the data in the following table were obtained. The data should fit a linear model of the form log k =
> For a Pt │Ce4+, Ce3+ half-cell, find the potential for the same ratios of [Ce4+]/[Ce3+] as given in Problem 16-27 for [Fe3+]/[Fe2+].
> Explain the difference between a. A colloidal and a crystalline precipitate. b. A gravimetric precipitation method and a gravimetric volatilization method. c. Precipitation and coprecipitation. d. Peptization and coagulation of a colloid. e. Occlusi
> For a Pt |Fe3+,Fe2+ half-cell, find the potential for the following ratios of [Fe3+]/[ Fe2+]. 0.001, 0.0025, 0.005, 0.0075, 0.010, 0.025, 0.050, 0.075, 0.100, 0.250, 0.500, 0.750, 1.00, 1.250, 1.50, 1.75, 2.50, 5.00, 10.00, 25.00, 75.00, and 100.00
> Calculate 0 E for the process //
> Given the formation constants
> Compute 0 E for the process
> The solubility product for
> The solubility-product constant for
> The solubility-product constant for The solubility-product constant for Ni2P2O7 is 1.7 3 10213. Calculate E0 for the process
> The solubility-product constant for Ag2SO3 is 1.5 3 10214. Calculate E0 for the process
> The following half-cells are on the left and coupled with the standard hydrogen electrode on the right to form a galvanic cell. Calculate the cell potential. Indicate which electrode would be the cathode if each cell were short-circuited. a. Cu uCu211
> The data in the following table represent electrode potential E versus concentration c. a. Transform the data to E versus 2log c values. b. Plot E versus 2log c, and find the least-squares estimate of the slope and intercept. Write the least squares equa
> Define a. Digestion. b. Adsorption. c. Reprecipitation. d. Precipitation from homogeneous solution. e. Counter-ion layer. f. Mother liquor. g. Supersaturation.
> What mass of Cu1IO322 can be formed from 0.475 g of CuSO4 5H2O?
> Define buffer capacity.
> If the following half-cells are the right-hand electrode in a galvanic cell with a standard hydrogen electrode on the left, calculate the cell potential. If the cell were shorted, indicate whether the electrodes shown would act as an anode or a cathode.
> calculate the potential of a platinum electrode immersed in a solution that is a. 0.0513 M in K4Fe1CN26 and 0.00589 M in K3Fe1CN26. b. 0.0300 M in FeSO4 and 0.00825 M in Fe21SO423. c. buffered to a pH of 4.85 and saturated with H2 at 1.00 atm. d. 0
> calculate the potential of a platinum electrode immersed in a solution that is a. 0.0160 M in K2PtCl4 and 0.2450 M in KCl. b. 0.0650 M in Sn1SO422 and 3.5 3 1023 M in SnSO4. c. buffered to a pH of 6.50 and saturated with H21g2 at 1.00 atm. d. 0.0
> Use activities to calculate the electrode potential of a hydrogen electrode in which the electrolyte is 0.0200 M HCl and the activity of H2 is 1.00 atm.
> Calculate the potential of a zinc electrode immersed in
> Calculate the potential of a copper electrode immersed in
> Consider the following oxidation/reduction reactions: a. Write each net process in terms of two balanced half-reactions. b. Express each half-reaction as a reduction. c. Arrange the half-reactions in (b) in order of decreasing effectiveness as electron
> Consider the following oxidation/reduction reactions: a. Write each net process in terms of two balanced half-reactions. b. Express each half-reaction as a reduction. c. Arrange the half-reactions in (b) in order of decreasing effectiveness as electron
> Identify the oxidizing agent and the reducing agent on the left side of each equation in Problem 16-9; write a balanced equation for each half-reaction.
> Calculate the solubility of the solutes in Problem 7-10 for solutions in which the anion concentration is 0.030 M.
> Generate the solubility-product expression for a. CuBr. b. MgCO3. C. PbCl2. d. CaSO4. e. Ag3AsO4.
> Write balanced net ionic equations for the following reactions. Supply H+ and/or H O2 as needed to obtain balance.
> Average human blood contains 300 nmoles of hemoglobin(Hb) per liter of plasma and 2.2 mmol per liter of whole blood. Calculate a. The molar concentration in each of these media. b. pHb in plasma in human serum.
> Identify the oxidizing agent and the reducing agent on the left side of each equation in Problem 16-7; write a balanced equation for each half-reaction.
> Write balanced net ionic equations for the following reactions. Supply H+ and/or H O2 as needed to obtain balance.
> The standard electrode potential for the reduction of Ni2+ to Ni is −. 0 25 V. Would the potential of a nickel electrode immersed in a 1.00 M NaOH solution saturated with Ni(OH 2) be more negative than E 0Ni2+ /Ni or less? Explain.
> In what respect is the Fajans method superior to the Volhard method for the titration of chloride ion?
> Write chemical formulas for the following complex ions:
> Explain how stepwise and overall formation constants are related.
> Write chemical equations and equilibrium-constant expressions for the stepwise formation of
> When a 100.0-mL portion of a solution containing 0.500 g of AgNO3 is mixed with 100.0 mL of a solution containing 0.300 g of K2CrO4, a bright red precipitate of Ag2CrO4 forms. a. Assuming that the solubility of Ag2CrO4 is negligible, calculate the mass
> Describe three general methods for performing EDTA titrations. What are the advantages of each?
> The data in the following table were obtained during a colorimetric determination of glucose in blood serum. a. Assuming a linear relationship between the variables, find the least-squares estimates of the slope and intercept. b. What are the standard de
