39. Sketch a voltaic cell for each redox reaction. Label the anode and cathode and indicate the half-reaction that occurs at each electrode and the species present in each solution. Also indicate the direction of electron flow. a. 2 Ag+ (aq) + Pb (s) → 2 Ag (s) + Pb²+ (aq) b. 2 C1O₂ (g) +2 I (aq) → 2 ClO₂¯¯ (aq) + I₂ (s) c. O₂ (g) + 4 H+ (aq) + 2 Zn (s) → 2 H₂O (1) + 2 Zn²+ (aq) 2+

39. Sketch a voltaic cell for each redox reaction. Label the anode and cathode and indicate the half-reaction that occurs at each electrode and the species present in each solution. Also indicate the direction of electron flow. a. 2 Ag+ (aq) + Pb (s) → 2 Ag (s) + Pb²+ (aq) b. 2 C1O₂ (g) +2 I (aq) → 2 ClO₂¯¯ (aq) + I₂ (s) c. O₂ (g) + 4 H+ (aq) + 2 Zn (s) → 2 H₂O (1) + 2 Zn²+ (aq) 2+

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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Suppose a galvanic cell is established using the following cell reaction. Cu (s) + 2Ag^1 (0.100 M) <----> Cu^2+ (0.200 M) + 2Ag (s) Each compartment of the cell contains 125 mL of solution.1. By how many grams will the total mass of the two electrodes change during this time period?a. The change in mass of the copper electrode =b. The change in mass of the silver electrode = c. The total change in mass of the two electrodes = 2. How will the entire mass of the galvanic cell be affected?
1. Draw the cell diagrams for each of the redox reactions given below. You may use platinum as an inert electrode. a. 2 Al(s) + 3 H₂SO4(aq) → Al₂(SO4)3(aq) + 3 H₂(g) b. Fe₂(SO4)3 + 3 Pb(s) → 3 PbSO4(s) + 2 Fe(s) c. CuSO4(aq) + H₂(g) → Cu(s) + H₂SO4(aq) d. 2 NaBr(aq) + 12(g) → Br₂() +2 Nal(aq) e. 2 FeCl₂(aq) + SnCl₂(aq) → 2 FeCl3(aq) + Sn(s) f. 2 FeCl₂(aq) + SnCl(aq) → 2 FeCl3(aq) + SnCl₂(aq) See prob 18.1 for 2. Consider a galvanic cell constructed from the following half-cells that are linked by a porous membrane. (1) an Au electrode dipped into 1.0 M Au(NO3)3 and (2) an Fe electrode dipped into 1.0 M FeSO4. Answer the following questions: a. Which electrode is the cathode? b. Write a balanced equation for the reaction occurring while the cell is discharging. c. What emf should the cell generate? Use Fig 18.1 for d. In what direction will electrons flow in the outer circuit? e. Toward which electrode will positive ions migrate?
3. In an electrochemical cell oxidation and reduction reactions occurs simultaneously, and it is these same reactions which generates electrical energy and powers the cell. a) Clearly identify the difference between an Oxidation and a Reduction reaction. b) Given the overall equation for the cell reaction: 2Cu*(aq) + I2 (aq) → 2Cu²*(aq) + 2I(aq) i. Identify the species which are the oxidizing agent and the reducing agent. Explain your answers. ii. Write the two half equations. iii. Draw the cell diagram iv. Calculate the standard cell e.m.f. v. Suggest whether the cell reaction is feasible. Explain your answer.
Find the voltage of the cell if the right half-cell contains 0.330 M AgNO₃ (aq) and the left half-cell contains 0.115 M Cu(NO₃)₂ (aq). Express your answer to three significant figures.
An unknown metal, X, has been deposited on the cathode of an electrolytic cell using an electric current of 3.30 Amps for a period of 25.0 mins: 2X2+ + 2 e– → X The mass of X collected at the cathode was 1.51 g. Identify the unknown metal, X.
4. For the following redox reaction in a galvanic cell, write the oxidation half-reaction and the reduction-half reaction, and calculate the standard cell potential of the reaction. Use Table 1 from the background as needed. Explain how you identified which substance was oxidized and which was reduced. Show all of your work. Redox reaction: Cu(s) + 2 Fe* (aq) → Cu?+ (aq) + 2 Fe* (aq) Table 1. Standard reduction potentials. Half-Reaction E°(Volts) F,(g) + 2e → 2F (aq) CI, (g) + 2e → 2CI(aq) Br, () + 2e → 2Br(aq) Ag*(aq) + e¯ → Ag(s) +2.87 +1.36 +1.07 +0.80 Fe*(aq) + e → Fe²"(aq) +0.77 Cu²*(aq) + 2e' → Cu(s) +0.34 2H*(aq) + 2e¯ → H,(g) 0.00 Fe2*(aq) + 2e' → Fe(s) -0.44 Zn*(aq) + 2e' → Zn(s) -0.76 Al³*(aq) + 3e' → Al(s) -1.66 Mg"(aq) + 2e' → Mg(s) -2.37 Ca²"(aq) + 2e" → Ca(s) -2.87 K*(aq) + e' → K(s) -2.93 В U T T O Word(s)
When an aqueous NaCl solution is electrolyzed, how many coulombs of electric charge need to be transferred at the anode to release 0.130 molmol of Cl2Cl2 gas? 2Cl−(aq)→Cl2(g)+2e−2Cl−(aq)→Cl2(g)+2e−
For each of the reactions, label the half-reaction that is the anode and the reaction that is the cathode in a Galvanic CellI. 1. Cu(s) + 2 Ag*(aq) → Cu²+(aq) + 2 Ag(s) → Zn2+(aq) + H2(g) 2. Zn(s) + 2H*(aq) → Cu(s) → Cu2+(aq) + 2e Choose... + Ag*(aq) + 1e → Ag(s) Choose... + Zn(s) Zn(s) → Zn2+(aq) + 2e Choose... + 2H*(aq) + 2e → H2(g) Choose... + Don Choose... anode cathode
A galvanic cell at a temperature of 25.0 °C is powered by the following redox reaction: Cu?* (aq)+Zn (s) → Cu (s) +Zn²* (aq) Suppose the cell is prepared with 3.40 M Cu´" in one half-cell and 5.47 M Zn2+ in the other. Calculate the cell voltage under these conditions. Round your answer to 3 significant digits. x10
What is the shorthand notation that represents the following galvanic cell reaction? 2 Fe2+ (aq) + Cl2(g) → 2 Fe3+(aq) + 2 Cl-(aq)
1. Mercuric oxide dry-cell batteries are often used where a high-energy density is required, such as in watches and cameras. The two half-cell reactions that occur in the battery are HgO(s)+H2O(l)+2e−→Hg(l)+2OH−(aq) Zn(s)+2OH−(aq)→ZnO(s)+H2O(l)+2e− (a) Write the overall cell reaction. (b) The value of E∘redfor the cathode reaction is +0.098V . The overall cell potential is +1.35V . Assuming that both half-cells operate under standard conditions, what is the standard reduction potential for the anode reaction?
1. Balance the following Redox Reactions. Always include the state of each species: (s), (I), (g), or (aq). a. Zn(s) + NO3- (aq) → Zn(OH)4^2¯(aq) + NH3(1) (basic) b. CHC13(1) + MnO4 - (aq) → Cl2(g) + CO2(g) + Mn2+(aq) (acidic) (Note: Cl here is chlorine)