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.
> 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?
> 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.
> 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
> Why are multidentate ligands preferable to unidentate ligands for complexometric titrations?
> Define a. ligand. b. chelate. c. tetradentate chelating agent. d. adsorption indicator. e. argentometric titration. f. conditional formation constant. g. EDTA displacement titration. h. water hardness.
> What is a buffer solution, and what are its properties?
> A 50.0-mL portion of a solution containing 0.200 g of BaCl2 . 2H2O is mixed with 50.0 mL of a solution containing 0.300 g of NaIO3. Assume that the solubility of Ba1IO322 in water is negligibly small and calculate a. The mass of the precipitated Ba1IO32
> The following are relative peak areas for chromatograms of standard solutions of methyl vinyl ketone (MVK). a. Determine the coefficients of the best fit line using the least-squares method. b. Construct an ANOVA table. c. Plot the least-squares line as
> Seawater contains an average of 1.08 × 103 ppm of Na+ and 270 ppm of SO42-. Calculate a. The molar concentrations of Na+ and SO2- given that the average density of seawater is 1.02 g/mL. b. The pNa and pSO4 for seawater.
> A 6.881-g sample containing magnesium chloride and sodium chloride was dissolved in sufficient water to give 500 mL of solution. Determination of the chloride content of a 50.0-mL aliquot resulted in the formation of 0.5923 g of AgCl. The magnesium in a
> The following data were obtained in calibrating a calcium ion electrode for the determination of pCa. A linear relationship between the potential and pCa is known to exist. a. Plot the data, and draw a line through the points by eye. b. Find the least-s
> At 25°C, what is the hydronium ion concentration in a. 0.100 M chloroacetic acid? b. 0.100 M sodium chloroacetate? c. 0.0300 M methylamine? d. 0.0300 M methylamine hydrochloride? e. 1.50 3 1023 M aniline hydrochloride? f. 0.200 M HIO3?
> What mass in grams of CO2 is evolved in the complete decomposition of a 2.300-g sample that is 38.0% MgCO3 and 42.0% K2CO3 by mass?
> A solution contains 1.569 mg of CoSO4 (155.0 g/mol) per milliliter. Calculate
> A solution was prepared by dissolving about 3.0 g of Na2H2Y⋅2H2O in approximately 1 L of water and standardizing against 50.00-mL aliquots of 0. 00397MMg2+ An average titration of 30.27 mL was required. Calculate the molar concentration of the EDTA.
> An EDTA solution was prepared by dissolving 3.426 g of purified and dried Na2H2Y · 2H2O in sufficient water to give 1.000 L. Calculate the molar concentration, given that the solute contained 0.3% excess moisture (see Section 15D-1).
> Why is a small amount of MgY2− often added to a water specimen that is to be titrated for hardness?
> The sulfate ion concentration in natural water can be determined by measuring the turbidity that results when an excess of BaCl2 is added to a measured quantity of the sample. A turbidimeter, the instrument used for this analysis, was calibrated with a s
> Given an overall complex formation reaction of with an overall formation constant of βn, show that the following relationship holds:
> Propose a complexometric method for the determination of the individual components in a solution containing
> Write a conditional overall formation constant for [Fe(Ox)3] 3− in terms of α2 for oxalic acid and the β value for the complex. Also express the conditional constant in terms of concentrations as in Equation 15-20.
> Write conditional formation constants for 1:1 complexes of Al(III) with each of the ligands in Problem 15-11. Express these constants in terms of the α value and the formation constant and in terms of concentrations as in Equation 15-20.
> Write equations in terms of the acid dissociation constants and [H+] for the highest alpha value for each of the following weak acid ligands:
> The phosphorus in a 0.3019-g sample was precipitated as the slightly soluble 1NH423PO4 . 12MoO3. This precipitate was filtered, washed, and then redissolved in acid. Treatment of the resulting solution with an excess of Pb21 resulted in the formation of
> Convert the following p-functions to molar concentrations: a. pH = 1.102. b. pOH = 0.0057. c. pBr = 7.77. d. pCa = -0.221. e. pLi = 12.35. f. pNO3 = 0.054. g. pMn = 0.135. h. pCl = 8.92.
> What is the difference between molar species concentration and molar analytical concentration?
> Outline a method for the determination of K+ based on argentometry. Write balanced equations for the chemical reactions.
> Outline a method for the determination of K+ based on argentometry. Write balanced equations for the chemical reactions.
> Why does the charge on the surface of precipitate particles change sign at the equivalence point of a titration?
> A method for the determination of the corticosteroid methylprednisolone acetate in solutions obtained from pharmaceutical preparations yielded a mean value of 3.7 mg mL-1with a standard deviation of 0.3 mg mL-1. For quality control purposes, the relative
> Briefly explain why the sparingly soluble product must be removed by filtration before you back-titrate the excess silver ion in the Volhard determination of
> Give two reasons why ( ( 3 2 KH IO is preferred over benzoic acid as a primary standard for a 0.010 M NaOH solution.
> Why is it common practice to boil the solution near the equivalence point in the standardization of Na2CO3 with acid?
> The boiling points of HCl and CO2 are nearly the same (285°C and 278°C). Explain why CO2 can be removed from an aqueous solution by boiling briefly while essentially no HCl is lost even after boiling for 1 hour or more.
> Four analysts perform replicate sets of Hg determinations on the same analytical sample. The results in ppb Hg are shown in the following table: a. State the appropriate hypotheses. b. Do the analysts differ at the 95% confidence level? At the 99% confi
> Describe how Na2CO3 of primary-standard grade can be prepared from primary-standard NaHCO3.
> Why is nitric acid seldom used to prepare standard acid solutions?
> A 10.00-mL sample of vinegar (acetic acid, CH3COOH) was pipetted into a flask, two drops of phenolphthalein indicator were added, and the acid was titrated with 0.1008 M NaOH a. If 45.62 mL of the base was required for the titration, what was the molar c
> At 25°C, what are the molar H3O1 and OH2 concentrations in a. 0.0300 M HCOOH? b. 0.0600 M HN3? c. 0.200 M ethylamine? d. 0.100 M trimethylamine? e. 0.250 M C6H5COONa (sodium benzoate)? f. 0.0750 M CH3CH2COONa? g. 0.250 M hydroxylamine hydrochloride? h. 0
> Calculate the equivalent mass of oxalic acid dehydrateH2C2O4 2H2O,12066g/mol when it is titrated to (a) a bromocresol green end point and (b) a phenolphthalein end point.
> Define the equivalent mass of (a) an acid and (b) a base
> The seller of a mining claim took a random ore sample that weighed approximately 5 lb and had an average particle diameter of 5.0 mm. Inspection revealed that about 1% of the sample was argentite (see Problem 6-9), and the remainder had a density of abou
> How did the definition of the mole change with the 2019 redefinition of SI base units?
> A series of solutions containing NaOH,Na2CO3 , and NaHCO3 , alone or in compatible combination, was titrated with 0.1202 M HCl. In the following table are the volumes of acid needed to titrate 25.00-mL portions of each solution to (1) a phenolphthalein a
> A series of solutions containing NaOH,Na3AsO4 , and Na2HAsO4 , alone or in compatible combination, was titrated with 0.08601 M HCl. In the following table are the volumes of acid needed to titrate 25.00-mL portions of each solution to (1) a phenolphthale
