[3] ~ Electrolytes Consider the following half-cells, where you are given their reduction potentials and molalities: Eºr red=+0.80 V & red=0.31 V * Ag(s) Agt (m1), NO3- (m1); * Pb(s) PbSO4(s), SO4²- (m2), K+ (2m2); where m₁ = 0.003 mol/kg and m2-0.009 mol/kg. (a) Make a scheme of a Galvanic battery that involves these half cells and show which are the anode and cathode. (Make a reasonable choice of salt bridge that reduces the liquid juncture potential AOLJ.) (b) Write the Nernst equation, using the Debye-Hückel limit law for the activity coefficients. (c) Calculate the electromotive force for the battery at 298K.

[3] ~ Electrolytes Consider the following half-cells, where you are given their reduction potentials and molalities: Eºr red=+0.80 V & red=0.31 V * Ag(s) Agt (m1), NO3- (m1); * Pb(s) PbSO4(s), SO4²- (m2), K+ (2m2); where m₁ = 0.003 mol/kg and m2-0.009 mol/kg. (a) Make a scheme of a Galvanic battery that involves these half cells and show which are the anode and cathode. (Make a reasonable choice of salt bridge that reduces the liquid juncture potential AOLJ.) (b) Write the Nernst equation, using the Debye-Hückel limit law for the activity coefficients. (c) Calculate the electromotive force for the battery at 298K.

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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Consider the following half-cells: (1) AgCl(s) | Ag(s), Cl–(aq) E°red = +0.22 V (2) Ni2+(aq) | Ni(s) E°red = –0.25 V (3) Fe2+(aq) | Fe(s) E°red = –0.44 V (4) Fe3+(aq) | Fe2+(aq) E°red = +0.77 V Which one of the following combinations of these half-cells will produce a galvanic cell generating the highest voltage under standard conditions? A. (2) and (3) B. (3) and (4) C. (1) and (4) D. (1) and (2) E. (2) and (4)
For the cell shown, the measured cell potential, Ecell, is -0.3651 V at 25 °C. Pt(s) | H₂(g, 0.797 atm) | H+ (aq, ? M) || Cd²+ (aq, 1.00 M) | Cd(s) The balanced reduction half-reactions for the cell, and their respective standard reduction potential values, Eº, are 2 H+ (aq) + 2 e- → H₂(g) Cd²+ (aq) + 2e → Cd(s) Calculate the H+ concentration. [H+] = x10 TOOLS - Eº = 0.00 V Eº = -0.403 V M
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Using these metal ion/metal standard reduction potentials: Cd2+ (aq)| Cd(s): -0.38 V; Zn2+ (aq) Zn(s): -0.78 V; Ni2+ (aq) |Ni (s): -0.29 V; Cu2+ (aq)| Cu(s): 0.30 V. Calculate the standard cell potential in V for the cell reaction Cu²+ (aq) + Cd(s)→Cd²+ (aq) + Cu(s) Report your answer with 2 places past the decimal point. Do not put unit in your answer. Type your answer...
4) So far, we have looked at finding cell potentials under standard conditions, however in reality, we are seldom at standard conditions. Recall that our equation for the Gibbs free energy at nonstandard conditions was AG=AG+RTINQ By substituting -nFE for G in the above equation we come to the Nernst equation (shown below), which gives a way to predict E under any conditions, not just standard conditions. In this equation, R is the gas constant (8.314 J mol K). F is the Faraday Constant (96480 C mol) and Q is the reaction quotient. RT E=E no cell The Nemst equation is extremely powerful, and can be written in a slightly more useful form. First, the natural log (In) can be replaced by log base 10 (log) by multiplying by 2.3 (InQ = 2.3logQ), and we can rearrange terms to put the equation into the form of a straight line, as shown. El 2.3RT nFlogQ+ Ecell a) Consider this form of the Nernst equation, sketch a plot of E. vs logQ. b) Notice that the term "-(2.3*RT/nF)" is made up entirely of…
Consider the following half-reactions: Half-reaction E° (V) Ag*(aq) + e" Co2*(aq) + 2e → Ag(s) 0.799V → Co(s) |-0.280V Zn2*(aq) + 2e → Zn(s) -0.763V (1) The strongest oxidizing agent is: enter formula (2) The weakest oxidizing agent is: (3) The weakest reducing agent is: (4) The strongest reducing agent is: (5) Will Zn2*(aq) oxidize Ag(s) to Ag"(aq)? (6) Which species can be reduced by Co(s)? If none, leave box blank.
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Give the answer to the questions below: Is AuCl more soluble in pure water or in 1M aqueous solution of NaCN? Will the Pb²+ or Pb half-cell (Eºred = -0.126 V) be the cathode when coupled with the SHE in a galvanic cell? What is the oxidizing agent in the following reaction? Mg(s) | Mg2+ (aq) || Pb2+ (aq) | Pb(s).
2. (a) A voltaic cell is constructed as follows: Zn | Zn²+ | H*|H₂ The half-cell reactions are: Zn →Zn²+ze E = 0.76 V If the voltage of the cell is 0.45V at 25°C when [Zn²] = 1.0 M and PH₂ = 1.0 atm, what is [H+]? 2H+ + 2e- → H₂ EⓇ = 0.0V (b) For the voltaic cell described in part (a), find the pH of the cathodic solution when pH2 = 1.0 atm and [Zn²+]= 0.1 M and the E*cell = 0.542 V. (c) What amount of time is needed to plate 2.08 g of copper at a constant current flow of 1.26 A from a 0.1 M CUSO4 solution? Cu²+ +ze--- Cu(s)
For the cell shown, the measured cell potential, Ecell, is -0.3619 V at 25 °C. Pt(s) | H₂(g, 0.721 atm) | H¹ (aq, ? M) || Cd²+ (aq, 1.00 M) | Cd(s) The balanced reduction half-reactions for the cell, and their respective standard reduction potential values, Eº, are 2 H+ (aq) + 2 e H₂(g) Eº = 0.00 V Cd²+ (aq) + 2e → Cd(s) E° = -0.403 V Calculate the H+ concentration. [H+] = TOOLS x10 M
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+|| Σ D. For the cell shown, the measured cell potential, Ecell, is-0.3621 V at 25 °C. Pt(s) | H, (g, 0.853 atm) | H* (aq, ? M) || Cd²+(aq, 1.00 M) | Cd(s) The balanced reduction half-reactions for the cell, and their respective standard reduction potential values, E°, are 2H*(aq) + 2 e¯ – H, (g) A 00'0 = 07 24, Cd2+ (aq) + 2 e Cd(s) E° = -0.403 V %3D Calculate the Ht concentration. M. = [,H] 10:23 AM M. 11/9/2021 7734 F4 F5 F7 F8 F10 F11 F12 PrtScr Insert Delete Eゴ #3 Backspace 2$ ) 4. 5. 9. 7. R } 1 Enter F. K. H. Sh