1. Consider the following standard reduction potentials in acid solution: E°(V) -1.66 Al³+ + 3e → Al(s) AgBr(s)+e → Ag(s) + Br- +0.07 Sn4+ + 2e + 2e +0.14 Fe³+ + e +0.77 The strongest reducing agent among those shown above is Hint: If answer is with charge, type this way: X2+ → Sn²+ → Fe²+

1. Consider the following standard reduction potentials in acid solution: E°(V) -1.66 Al³+ + 3e → Al(s) AgBr(s)+e → Ag(s) + Br- +0.07 Sn4+ + 2e + 2e +0.14 Fe³+ + e +0.77 The strongest reducing agent among those shown above is Hint: If answer is with charge, type this way: X2+ → Sn²+ → Fe²+

Chemistry

10th Edition

ISBN:9781305957404

Author:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste

Publisher:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste

Chapter1: Chemical Foundations

Section: Chapter Questions

Problem 1RQ: Define and explain the differences between the following terms. a. law and theory b. theory and...

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1.A metal "X" is used as a “sacrificial” anode on hulls of salt water vessels made of iron. The effect is to reverse the oxidation of iron to iron (II) by causing X to oxidize to the X ion. Calculate the voltage (in V, to two decimal places) for this reaction if it takes place under standard conditions. The standard reduction potential for X is -0.52 V (minus 0.52 V). 2.Batteries are charged and discharged due to the concentration of the solutions inside. Calculate the voltage of a battery made from zinc in a 0.1M solution of Zn2+ and a metal X in a 1 M X2+ solution using the table of reduction potentials and Eq. 6 at 25 degrees C. Answer with the unit V, to two decimal places. The standard reduction potential of X is -0.46 volts.
A metal "X" is used as a “sacrificial” anode on hulls of salt water vessels made of iron. The effect is to reverse the oxidation of iron to iron (II) by causing X to oxidize to the X ion. Calculate the voltage (in V, to two decimal places) for this reaction if it takes place under standard conditions. The standard reduction potential for X is -1.09 V (minus 1.09 V).
Refer to the galvanic cell below (the contents of each half-cell are written beneath each compartment): Ni Ag 1.0 M Ni2+ 1.0 M Ag* The standard reduction potentials are as follow: Ni2+ + 2 e- → Ni(s), E° = -0.25 Ag* + e - Ag(s), E° = 0.7996 When current is allowed to flow, which species is oxidised? Select one: а. Ag b. Cannot be determined from the data given. C. Ni2+ O d. Ag* O e. Ni
A copper, Cu(s), electrode is immersed in a solution that is 1.00 M in ammonia, NH3, and 1.00 M in tetraamminecopper(II), [Cu(NH3)4²+. If a standard hydrogen electrode is used Use the standard reduction potentials shown here to answer the questions. as the cathode, the cell potential, Ecell is found to be 0.080 V at 298 K. Reduction half-reaction E° (V) Cu2(aq) + 2e¬Cu(s) 0.337 Part A 2H+(aq) + 20¬H2(g) 0.000 Based on the cell potential, what is the concentration of Cu2+in this solution? Express your answer to one significant figure and include the appropriate units. Templates Symbols uado redo reset keyboard shortcuts help [Cu² = Value Units
Consider the following standard reduction potentials in acid solution: E°(V) Al⁺³ + 3e− → Al(s) −1.66 Ag⁺¹ + e− → Ag(s) +0.80 Sn⁺⁴ + 2e− → Sn⁺² +0.14 Fe⁺³ + e− → Fe⁺² +0.77 The strongest reducing agent among those shown above is A) Al⁰ B) Fe⁺³ C) Br⁻ D) Al⁺³
Standard Reduction Potentials at 25 °C Half-reaction. Eº(volts) Au³+ + 2e Au+ 1.290 Aut+eAu 1.680 Au³+ + 3e →→Au 1.420 In a disproportionation reaction, the same species is oxidized and reduced. Use the data given above to analyze the disproportionation reaction: 3Aut 2Au + Au³+ Will Au* disproportionate in aqueous solution? To justify your answer, what is the standard cell voltage for the disproportionation reaction?
Based on standard reduction potentials (in second table below), will each of these reactions (in table 1) take place? How to calculate this?
Consider the following two half-reactions and their standard reduction potentials. N2(g) +H,0(1) E° = 1.766 V MnO, (s) + 2H,0(1) E° = 1.673 V N,0(g) + 2H*(aq) + 2e¯ → MnO,(aq) + 4H*(aq)+3e¯ 1st attempt l See Calculate E°for the spontaneous redox reaction that occurs when these two half-reactions are coupled. V
Enter electrons as e". Use smallest possible integer coefficients. States are not required. If a box is not needed, leave it blank. Standard Reduction Potentials (Volts) at 25 °C Cl₂ (g) + 2 e →2 Cl(aq) O2(g) + 4 H3O+ (aq) + 4 e 6 H₂O(1) 1.360 1.229 2 H2O(l) + 2e Al3+ (aq) + 3 e → Al (s) H2(g) + 2 OH(aq) -0.828 -1.660 A 0.114 M neutral aqueous AlCl3 solution is electrolyzed under 1 atm pressure using platinum electrodes. (a) Write the half-reactions expected to occur, taking into account the effect of the non-standard conditions on the electrode potentials. Half-reaction at anode: + + + + Half-reaction at cathode: + + + + (b) What is the expected decomposition potential?
2+ 4+ a. Will Fe+ oxidize Sn to Sn*t ? Assume that the reaction occurs at standard electrochemical conditions. Half-Reaction e° (V) Fes+ +e → Fe2+ +0.77 Sn n4+ + 2e → Sn²+ +0.15 O yes O no b. Will dichromate ions oxidize chloride ions to free chlorine in acidic solution? Assume that the reaction occurs at standard electrochemical conditions. Half-Reaction e° (V) Cr2 0, + 14H+ + 6e¯ → 2Cr³+ + 7H2O +1.33 Cl2 + 2e → 2C1 +1.36 O yes O no
Which one of the following reactions will occur spontaneously under standard conditions (and 25°C) ? You should consult a table of Standard Reduction Potentials 3Ag + A1³+→ 3Ag+ + Al 2A13+ + 3Fe → 2Al + 3Fe2+ Cu + 2H+ Cu²+ + H₂ 3Mg2+ + 2Cr→ 3Mg + 2Cr³+ Mg2+ + Ca → Mg + Ca²+
Use standard reduction potentials to calculate the standard free energy change in kJ for the reaction: 2H+ (aq) + Hg(1)→ H₂(g) + Hg²+ (aq) 2H+ (aq) + 2e¯ → H₂(g) Fº red 2+ Hg2 (aq) + 2e¯ → Hg(1) AGO = || kJ K for this reaction would be Eº red = 0.000 V = = 0.855 V than one.