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Quantitative Chemical Analysis 8th edition Daniel C. Harris - Solutions

How many milliliters of 0.0500 M EDTA are required to react with 50.0 mL of 0.0100 M Ca2+? With 50.0 mL of 0.0100 M Al3+?
A 50.0-mL sample containing Ni2+ was treated with 25.0 mL of 0.050 0 M EDTA to complex all the Ni2+ and leave excess EDTA in solution. The excess EDTA was then back titrated, requiring 5.00 mL of 0.0500 M Zn2+. What was the concentration of Ni2+ in the original solution?
A 50.0-mL aliquot of solution containing 0.450 g of MgSO4 (FM 120.37) in 0.500 L required 37.6 mL of EDTA solution for titration. How many milligrams of CaCO3 (FM 100.09) will react with 1.00 mL of this EDTA solution?
Cyanide solution (12.73 mL) was treated with 25.00 mL of Ni2+ solution (containing excess Ni2+) to convert the cyanide into tetracyanonickelate(II): 4CN- + Ni2+ → Ni(CN)2-4 Excess Ni2+ was then titrated with 10.15 mL of 0.01307 M EDTA. Ni(CN)42- does not react with EDTA. If 39.35 mL of EDTA were
A 1.000-mL sample of unknown containing Co2+ and Ni2+ was treated with 25.00 mL of 0.03872 M EDTA. Back titration with 0.02127 M Zn2+ at pH 5 required 23.54 mL to reach the xylenol orange end point. A 2.000-mL sample of unknown was passed through an ion-exchange column that retards CO2+ more than
A 50.0-mL solution containing Ni2+ and Zn2+ was treated with 25.0 mL of 0.0452 M EDTA to bind all the metal. The excess unreacted EDTA required 12.4 mL of 0.012 3 M Mg2+ for complete reaction. An excess of the reagent 2,3-dimercapto-1-propanol was then added to displace the EDTA from zinc. Another
Sulfide ion was determined by indirect titration with EDTA. To a solution containing 25.00 mL of 0.04332 M Cu(ClO4)2 plus 15 mL of 1M acetate buffer (pH 4.5) were added 25.00 mL of unknown sulfide solution with vigorous stirring. The CuS precipitate was filtered and washed with hot water. Then
Cesium ion does not form a strong EDTA complex, but it can be analyzed by adding a known excess of NaBiI4 in cold concentrated acetic acid containing excess NaI. Solid Cs3Bi2I9 is precipitated, filtered, and removed. The excess yellow BiI4- is then titrated with EDTA. The end point occurs when the
The sulfur content of insoluble sulfides that do not readily dissolve in acid can be measured by oxidation with Br2 to SO2-4.25 Metal ions are then replaced with H+ by an ion exchange column, and sulfate is precipitated as BaSO4 with a known excess of BaCl2. The excess Ba2+ is then titrated with
By analogy to a hydrogen ion buffer, a metal ion buffer tends to maintain a particular metal ion concentration in solution. A mixture of the acid HA and its conjugate base A- maintains [H+] defined by the equation Ka = [A-][H+]/[HA]. A mixture of CaY2- and Y4- serves as a Ca2- buffer governed by
To measure oxygen isotopes in SO24- for geologic studies, SO24- was precipitated with excess Ba2+.24 In the presence of HNO3, BaSO4 precipitate is contaminated by NO3-. The solid can be purified by washing, redissolving in the absence of HNO3, and reprecipitating. For purification, 30 mg of BaSO4
A 100.0 mL solution of 0.0500 M Mn buffered to pH 9.00 was titrated with 0.0500 M EDTA.(a) What is the equivalence volume, Ve, in milliliters?(b) Calculate the concentration of Mn+ at V = 1/2 Ve+.(c) What fraction (α γ4-) of free EDTA is in the form γ 4- at pH 9.00?(d) The formation constant
Calculate pCo2+ at each of the following points in the titration of 25.00 mL of 0.020 26 M Co2+ by 0.03855 M EDTA at pH 6.00: (a) 12.00 mL; (b) Ve; (c) 14.00 mL.
Consider the titration of 25.0 mL of 0.0200 M MnSO4 with 0.010 0 M EDTA in a solution buffered to pH 8.00. Calculate pMn2+ at the following volumes of added EDTA and sketch the titration curve:(a) 0 mL (b) 20.0 mL (c) 40.0 mL (d) 49.0 mL(e) 49.9 mL(f) 50.0 mL (g) 50.1 mL(h) 55.0 mL(i) 60.0 mL
For the same volumes used in Problem 11-8, calculate pCa2+ for the titration of 25.00 mL of 0.02000 M EDTA with 0.010 00 M CaSO4 at pH 10.00.
Why does the solubility of a salt of a basic anion increase with decreasing pH? Write chemical reactions for the minerals galena (PbS) and cerussite (PbCO3) to explain how acid rain mobilizes traces of metal from relatively inert forms into the environment, where the metals can be taken up by
Considering just acid-base chemistry, not ion pairing and not activity coefficients, find the pH and composition of 1.00 L of solution containing 0.040 mol H4EDTA (EDTA) = ethylenedinitrilotetraacetic acid = H4O , 0.030 mol lysine (molecule) = , and 0.050 mol NaOH.
The solution containing no added KNO3 for Figure 7-1 contains 5.0 mM Fe(NO3)3, 5.0 M NaSCN, and 15 mM HNO3. We will use Davies activity coefficients to find the concentrations of all species in the following reactions:This problem is harder than you might think, so we will guide you through the
(a) Follow the steps of Problem 12-11 to solve this one. From the following equilibria, find the concentrations of species and the pH of 1.0 mM La2(SO4)3. Use Davies activity coefficients. You may wish to use Goal Seek in place of Solver as described in Problem 12-11.
Find the pH and composition of saturated CaSO4 in water, considering Reactions 7-23 to 7 27 and using Davies activity coefficients. Suggested procedure: Use activity coefficients in all equilibrium expressions. Let the initial ionic strength be 0. (i) In the mass balance 7-29, express [CaOH+] in
Find the composition of a saturated solution of AgCN containing 0.10 M KCN adjusted to pH 12.00 with NaOH. Consider the following equilibria and use Davies activity coefficients.Suggested procedure: Let [CN-] be the master variable. We know [H+] from the pH, and we can find [Ag+] = Ksp /[CN-]. (i)
Consider the reactions of Fe2+ with the amino acid glycine:
Data for the glycine difference plot in Figure 12-14 are given below.
A solution containing 0.139 mmol of the triprotic acid tris(2-aminoethyl)amine ∙ 3HCl plus 0.115 mmol HCl in 40 mL of 0.10 M KCl was titrated with 0.490 5 M NaOH to measure acid dissociation constants.(a)Write expressions for the experimental mean fraction of protonation,H(measured), and the
(a) For the following reactions, prepare a diagram showing log(concentration) versus pH for all species in the pH range 2 to 12 for a solution made by dissolving 0.025 mol CuSO4 in 1.00 L. Equilibrium constants apply at μ = 0.1 M, which you shouldassume is constant. Do not use activity
This problem incorporates ion-pair equilibria 12-12 and 12-13 into the acid-base chemistry of Section 12-1. (a) From mass balance 12-15, derive Equation 12-16. (b) Substitute equilibrium expressions into mass balance 12-17 to derive an expression for [T2-] in terms of [H+], [Na+], and various
(a) Considering just acid-base chemistry, not ion pairing and not activity coefficients, use the systematic treatment of equilibrium to find the pH of 1.00 L of solution containing 0.010 0 mol hydroxybenzene (HA) and 0.0050 mol KOH.(b) What pH would you have predicted from your knowledge of Chapter
Repeat part (a) of Problem 12-2 with Davies activity coefficients. Remember that pH = - log([H+]). γH+).
From pK1 and pK2 for glycine at μ = 0 in Table 9-1, compute pK1 and pK2 that apply at μ = 0.1 M. Use the Davies equation for activity coefficients. Compare your answer with experimental values in cells B10 and B11 of Figure 12 - 13.
Considering just acid-base chemistry, not ion pairing and not activity coefficients, find the pH and concentrations of species in 1.00 L of solution containing 0.040 mol benzene-1,2,3- tricarboxylic acid (H3A), 0.030 mol imidazole (a neutral molecule, HB), and 0.035 mol NaOH.
Considering just acid-base chemistry, not ion pairing and not activity coefficients, find the pH and concentrations of species in 1.00 L of solution containing 0.020 mol arginine, 0.030 mol glutamic acid, and 0.005 mol KOH.
Solve Problem 12-7 by using Davies activity coefficients.
A solution containing 0.008 695 m KH2PO4 and 0.03043 m Na2HPO4 is a primary standard buffer with a stated pH of 7.413 at 25°C. Calculate the pH of this solution by using the systematic treatment of equilibrium with activity coefficients from (a) The Davies equation and (b) The extended
Explain the difference between electric charge (q, coulombs), electric current (I, amperes), and electric potential (E, volts).
Consider the rechargeable battery:Zn(s) | ZnCl2(aq) || Cl-(aq) | Cl2(l) | C(s)(a) Write reduction half-reactions for each electrode. From which electrode will electrons flow from the battery into a circuit if the electrode potentials are not too different from E values?(b) If the battery delivers
(a) Ideal formulas for the electrodes of the Li+-ion battery described in the chapter opener are C6Li (FM 79.01) and LiCoO2 (FM 97.87). When the battery operates, C6Li is consumed and LiCoO2 is formed. Write a half-reaction for each electrode, assuming that x ≠ 1 in the reaction . (In fact, in
(a) Cyanide ion causes E for Fe(III) to decrease:Which ion, Fe(III) or Fe(II), is stabilized more by complexing with CN-?(b) Using Appendix H, answer the same question when the ligand is phenanthroline instead of cyanide.
Write the Nernst equation for the following half-reaction find E when pH = 3.00 and PASH3 mbar. As(s) + 3H+ + 3e- ⇌ AsH3(g) E° = - 0.238 Arsine
(a) Write the line notation for the following cell.(b) Calculate the potential of each half-cell and the cell voltage, E. In which direction will electrons flow through the circuit? Write the spontaneous net cell reaction.(c) The left half-cell was loaded with 14.3 mL of Br2(l)(density = 3.12
A nickel-metal hydride rechargeable battery formerly used in laptop computers is based on the following chemistry:Cathode:Anode: The anode material, MH, is a transition metal hydride or rare earth alloy hydride. Explain why the voltage remains nearly constant during the entire discharge cycle.
Suppose that the concentrations of NaF and KCl were each 0.10 M in the cell(a) Using the half-reactions 2AgCI(s) + 2e- ‡Œ 2A , calculate the cell voltage. (b) By the reasoning in Figure 13-8, in which direction do electrons flow? (c) Now calculate the cell voltage by using the reactions
(a) How many electrons are in one coulomb? (b) How many coulombs are in one mole of charge?
The following cell was set up to measure the standard reduction potential of the Ag+ | Ag couple:Pt(s) Æ’ HCl(0.010 00 M), H2(g) || AgNO3 (0.010 00 M) Æ’ Ag(s)The temperature was 25oC (the standard condition) and atmospheric pressure was 751.0 Torr. Because the vapor pressure
Write a balanced chemical equation (in acidic solution) for the reaction represented by the question mark on the lower arrow.18 Calculate E° for the reaction.
What must be the relation between E°1 and E°2 if the species X+ is to disproportionate spontaneously under standard conditions to X3+ and X(s)? Write a balanced equation for the disproportionation.
Including activities, calculate the voltage of the cell Ni(s) | NiSO4(0.0020 M) || CuCl2(0.003 0 M)| Cu(s). Assume that the salts are completely dissociated (that is, neglect ion-pair formation). By the reasoning in Figure 13-8, in which direction do electrons flow?
For the reaction,CO + ½ O2 ⇌ CO2 ∆Go - 257 kJ per mole of CO at 298 K. Find E and the equilibrium constant for the reaction.
Calculate E°, ∆G , and K for the following reactions. (a) 4Co3+ + 2H2O ⇌ 4Co2+ O2(g) + 4H+ (b) Ag(S2O3)32- + Fe(CN)4-6 ⇌ Ag(s) + 2S2O2- 3 + Fe(CN)3-6
A solution contains 0.100 M Ce3+, 1.00 × 10 -4 M Ce4+, 1.00 × 10 -4 M Mn2+, 0.100 M , and 1.00 M HClO4.(a) Write a balanced net reaction that can occur between species in this solution.(b) Calculate ∆G° and K for the reaction.(c) Calculate E for the conditions given.(d) Calculate ∆G° for
For the cell Pt(s) | VO2+(0.116 M), V3+(0.116 M), H+(1.57 M) || Sn2+(0.031 8 M), Sn4+(0.031 8 M) | Pt(s), E (not E ) = - 0.289 V. Write the net cell reaction and calculate its equilibrium constant. Do not use E values from Appendix H to answer this question.
Calculate E° for the half-reaction pd(OH)2(s) +2e- ⇌ pd(s) 2OH- given that Ksp for Pd(OH)2 is 3 × 10-28 and E° = 0.915 V for the reaction Pd2+ + 2e- ⇌ Pd(s).
From the standard potentials for reduction of Br2(aq) and Br2(l) in Appendix H, calculate the solubility of Br2 in water at 25°C. Express your answer as g/L.
The basal rate of consumption of O2 by a 70-kg human is about 16 mol of O2 per day. This O2 oxidizes food and is reduced to H2O, providing energy for the organism:O2 + 4H+ + 4e- ⇌ 2H2O(a) To what current (in amperes = C/s) does this respiration rate correspond? (Current is defined by the flow of
Given the following information, calculate the standard potential for the reaction, FeY- + e- ‡Œ FeY2- where Y is EDTA.
For modest temperature excursions away from 25 C, the change in E for a half-reaction can be written in the formwhere E°(T) is the standard reduction potential at temperature T (°C), and ˆ†T is (T - 25). For the reaction ,Al3+ + 3e- ‡Œ Al(s) dE° /dT = 0.533 mV/K
This problem is slightly tricky. Calculate E° , ˆ†G° , and K for the reaction.which is the sum of three half-reactions listed in Appendix H. Use ˆ†G° (= - nFE° ) for each of the half-reactions to find ˆ†G° for the net reaction. Note that, if you
Thermodynamics of a solid-state reaction. The following electrochemical cell is reversible at 1 000 K in an atmosphere of flowing O2(g):19(a) Write a Nernst equation for each half-cell. Write the net reaction and its Nernst equation. The activity of O2(g) is the same on both sides, and the activity
With Figure 13-10 as an example, explain what we mean when we say that there is equilibrium within each half-cell but not necessarily between the two half-cells.
The cell Pt(s) | H2(g, 1.00 bar) | H(aq, pH 3.60) || Cl-(aq, x M) | AgCl(s) | Ag(s) can be used as a probe to find the concentration of Cl in the right compartment. (a) Write reactions for each half-cell, a balanced net cell reaction, and the Nernst equation for the net cell reaction. (b) Given
The quinhy drone electrode was introduced in 1921 as a means of measuring pH.20Pt(s) ƒ 1:1 mole ratio of quinone(aq) and hydroquinone(aq),unknown pH ||Cl-(aq, 0.50 M) ƒ Hg2Cl2(s) ƒ Hg(l) ƒ Pt(s)
The voltage for the following cell is 0.490 V. Find Kb for the organic base RNH2. Pt(s) ƒ H2(1.00 bar) ƒ RNH2(aq, 0.10 M), RNH+3 Cl(aq, 0.050 M) ||' S.H.E.
The voltage of the cell shown here is - 0.246 V. The right half-cell contains the metal ion, M2+, whose standard reduction potential is - 0.266 V.Mc2+ + 2e - ‡Œ E° = - 0.266 VCalculate Kf for the metal-EDTA complex
The following cell was constructed to find the difference in Ksp between two naturally occurring forms of CaCO3(s), called calcite and aragonite.21buffer(pH 7.00) ƒ CaCO3(s, aragonite) ƒ PbCO3(s) ƒ Pb(s)Pb(s) ƒ PbCO3(s) ƒ CaCO3(s, calcite) ƒ buffer(pH 7.00)'
A 6.00-V battery is connected across a 2.00-k Ω resistor. (a) How many electrons per second flow through the circuit? (b) How many joules of heat are produced for each electron? (c) If the circuit operates for 30.0 min, how many moles of electrons will have flowed through the resistor? (d) What
Do not ignore activity coefficients in this problem. If the voltage for the following cell is 0.512 V, find Ksp for Cu(IO3)2. Neglect any ion pairing. Ni(s)|| NiSO4(0.0025 M) ' KIO3(0.10 M) | Cu(IO3)2(s) |Cu(s)
Explain what E° is and why it is preferred over E° in biochemistry.
We are going to find E ° for the reaction .(a) Write the Nernst equation for the half-reaction, using E from Appendix H.(b) Rearrange the Nernst equation to the form(c) The quantity (E° other terms) is E°. Evaluate E° for pH = 7.00.
Evaluate E° for the half-reaction (CN)2(g) + 2H+ + 2e- ⇌ 2HCN(aq)
Calculate E° for the reaction
HOx is a monoprotic acid with Ka = 1.4 × 10-5 and H2Red- is a diprotic acid with K1 = 3.6 10-4 and K2 = 8.1 × 10 -8. Find E for the reaction. HOx + e- ⇌ H2Red- E = 0.062 V
Given the following information, find Ka for nitrous acid, HNO2. NO-3 + 3H+ + 2e- ⇌ HNO2 + H2O Eo = 0.940 V E° = 0.433 V
Using the reaction HPO24- + 2H+ + 2e- ⇌ HPO2-3 = H2O E° = - 0.234 V and acid dissociation constants from Appendix G, calculate E for the reaction H2PO4- + H+ + 2e- ⇌ HPO2-3 + H2O
This problem requires Beer's law from Chapter 17. The oxidized form (Ox) of a flavor protein that functions as a one-electron reducing agent has a molar absorptivity ( ) of 1.12 × 104 M-1 cm -1 at 457 nm at pH 7.00. For the reduced form (Red), ε = 3.82 × 103 M-1cm-1 at 457 nm at pH 7.00. Ox + e-
Consider the redox reaction(a) Identify the oxidizing agent on the left side of the reaction and write a balanced oxidation half-reaction.(b) Identify the reducing agent of the left side of the reaction and write a balanced reduction half-reaction.(c) How many coulombs of charge are passed from
The space shuttle's expendable booster engines derive their power from solid reactants:6NH4+ ClO-4 (s) + 10Al(s) → 3N2(g) + 9H2O(g) + 5Al2O3(s) + 6HCI(g)(a) Find the oxidation numbers of the elements N, Cl, and Al in reactants and products. Which reactants act as reducing agents and which act as
Explain how a galvanic cell uses a spontaneous chemical reaction to generate electricity.
Write a line notation and two reduction half-reactions for each cell pictured above.
Draw a picture of the following cell and write reduction half- reactions for each electrode:
(a) Write the half-reactions for the silver-silver chloride and calomel reference electrodes.(b) Predict the voltage for the following cell.
Consider the cell,|| S.C.E cell solution |whose voltage is - 0.126 V. The cell solution contains 2.00 mmol of Fe(NH4)2(SO4)2, 1.00 mmol of FeCl3, 4.00 mmol of Na2EDTA, and lots of buffer, pH 6.78, in a volume of 1.00 L.(a) Write a reaction for the right half-cell.(b) Find the quotient [Fe2+]/[Fe3+]
Here's a cell you'll really like:The cell solution was made by mixing 25.0 mL of 4.00 mM KCN 25.0 mL of 4.00 mM KCu(CN)2 25.0 mL of 0.400 M acid, HA, with pKa = 9.50 25.0 mL of KOH solution The measured voltage was -0.440 V. Calculate the molarity of the KOH solution.
What causes a junction potential? How does this potential limit the accuracy of potentiometric analyses? Identify a cell in the illustrations in Section 13-2 that has no junction potential.
Why is the 0.1 M HCl | 0.1 M KCl junction potential of opposite sign and greater magnitude than the 0.1 M NaCl | 0.1 M KCl potential in Table 14-2?
Which side of the liquid junction 0.1 M KNO3 | 0.1 M NaCl will be negative?
Refer to the footnote in Table 14-1. How many seconds will it take for (a) H+ and (b) NO-3 to migrate a distance of 12.0 cm in a field of 7.80 × 103 V/m?
Suppose that an ideal hypothetical cell such as that inFigure 13-7 were set up to measure E° for the half-reaction .(a) Calculate the equilibrium constant for the net cell reaction.(b) If there were a junction potential of + 2 mV (increasing E from 0.799 to 0.801 V), by what percentage would the
Explain how the cell Ag(s) | AgCl(s) | 0.1 M HCl | 0.1 M KCl | AgCl(s) | Ag(s) can be used to measure the 0.1 M HCl | 0.1 M KCl junction potential.
The junction potential, Ej-, between solutions, and can be estimated with the Henderson equation:where zi is the charge of species i, ui is the mobility of species I (Table 14-1), Ci(α) is the concentration of species i in phase α , and Ci(β) is the concentration in phase . (Activity
Describe how you would calibrate a pH electrode and measure the pH of blood (~7.5) at 37°C. Use the standard buffers in Table 14-3.
List the sources of error associated with pH measurement with the glass electrode.
Why do glass pH electrodes tend to indicate a pH lower than the actual pH in strongly basic solution?
Suppose that the Ag | AgCl outer electrode in Figure 14-11 is filled with 0.1 M NaCl instead of saturated KCl. Suppose that the electrode is calibrated in a dilute buffer containing 0.1 M KCl at pH 6.54 at 25°C. The electrode is then dipped in a second buffer at the same pH and same temperature,
(a) When the difference in pH across the membrane of a glass electrode at 25°C is 4.63 pH units, how much voltage is generated by the pH gradient?(b) What would the voltage be for the same pH difference at 37°C?
When calibrating a glass electrode, 0.025 m potassium dihydrogen phosphate/0.025 m disodium hydrogen phosphate buffer (Table 14-3) gave a reading of - 18.3 mV at 20°C and 0.05 m potassium hydrogen phthalate buffer gave a reading of + 146.3 mV. What is the pH of an unknown giving a reading of +
The 0.025 0 m KH2PO4/0.025 0 m Na2HPO4 buffer (6) in Table 14-3 has a pH of 6.865 at 25°C.(a) Show that the ionic strength of the buffer is μ = 0.100 m.(b) From the pH and K2 for phosphoric acid, find the quotient of activity coefficients, , at = 0.100 m.(c) You have the urgent need to prepare
Explain the principle of operation of ion-selective electrodes. How does a compound electrode differ from a simple ion-selective electrode?
What does the selectivity coefficient tell us? Is it better to have a large or a small selectivity coefficient?
Suppose that the silver-silver chloride electrode in Figure 14-2 is replaced by a saturated calomel electrode. Calculate the cell voltageif [Fe2+] / [Fe3+] = 2.5 × 10-3.Figure 14-2
Why is it preferable to use a metal ion buffer to achieve pM = 8 rather than just dissolving enough M to give a 10-8 M solution?
To determine the concentration of a dilute analyte with an ion-selective electrode, why do we use standards with a constant, high concentration of an inert salt?

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