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13 Dec 2019

(Spontaneity of Redox Reactions – 13 points) For this entire question, assume T is constant at 300. K, H2(g) is maintained at standard state conditions, R = 8.314 J mol-1 K-1, F = 96485C mol-1, and all activity coefficients are equal to 1.0. a. (8 points) You have four beakers, each one containing 1.0 L of an aqueous solution buffered to a pH = 1.0 (i.e [H3O+] = [H+] = 0.1M) and with 1 M of dissolved metal ions in each. You place a bar of silver (Ag) in beaker #1 which contains 1M Ag+, a bar of copper in beaker #2 which contains 1M Cu2+, a bar of lead in beaker #3 which contains 1M Pb2+, and a bar of zinc in beaker #4 which contains 1M Zn2+. Using the table of standard reduction potentials provided to the right, identify and report which beaker(s) a spontaneous dissolution (M(s) ? Mn+(aq); where “M” represents the solid metal and “Mn+” represents the metal cations) of some portion of the metal bar would or would not occur – provide an answer for each of the four beakers and show some justification for your answer to receive any credit

b. (5 points) You now prepare a fifth beaker also containing 1.0 L of an aqueous solution buffered to a pH = 1.0 (i.e [H3O+] = [H+] = 0.1M). To this beaker you then add a bar of pure copper (Cu(s)). Will there be a spontaneous dissolution of some portion of the copper bar in order to the system to establishes equilibrium? If yes, utilize the provided table of standard reduction potentials and the Nernst equation to calculate and report the concentration of Cu2+ ions in solution once the system attains equilibrium. Note : this system is different than Part A above in which beaker #2 initially had 1M of Cu2+ ions present; in this beaker #5 there is initially zero Cu2+ ions. Report your final answer with 1 significant digit.

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Related questions

1. Consider a voltaic (galvanic) cell with the following metal electrodes. Identify which metal is the cathode and which is the anode, and calculate the cell potential.

(a) Al and Co(II)
Cathode: ____

Anode: ______

Ecell =

(b) Cd(II) and Ag(I)

cathode: ___

anode: ___

Ecell= ____

(c) Cr(III) and Sc(III)

cathode: ___

Anode:_____

Ecell:____

2. A voltaic cell contains two half-cells. One half-cell contains a titanium electrode immersed in a 1.00 M Ti(NO3)3 solution. The second half-cell contains a zinc electrode immersed in a 1.00 M Zn(NO3)2 solution.

Ti3+(aq) + 3 e− → Ti(s)

E⁰red  = −1.370 V
Zn2+(aq) + 2 e− → Zn(s)

E⁰red  = −0.762 V

(a) Using the standard reduction potentials given above, predict the standard cell potential of the voltaic cell.

(b) Write the overall balanced equation for the voltaic cell. (Include states-of-matter under the given conditions in your answer.)

3. ΔG° and E° can be said to measure the same thing, and are convertible by the equation

ΔG° = −nFE⁰cell 

where n is the total number of moles of electrons being transferred, and F is the Faraday constant 9.64853415 ✕ 104 C/mol. The free energy (ΔG°) of a spontaneous reaction is always negative.

For each of the electrochemical cells below, calculate the free energy of the system and state whether the reaction is spontaneous or non-spontaneous as written based on the cathode and anode assignment given. (Use the table of Standard Electrode Potentials.)

(a) The cathode is Zn(II) and the anode is Co(II).

free energy: ____ kJ

spontaneity: _____

1. (1 point) What are transferred in an oxidation-reduction reaction? *

• atoms

• electrons

• ions

• protons

2. (1 point) In the reaction of beryllium with fluorine, which atom is oxidized? *

• fluorine

• beryllium

• neither of them

• both of them

3. (1 point) Ag - - > Ag ^+ +1 e^- This equation represents the type of reaction called _____. *

• hydrolysis

• oxidation

• redox

• reduction

4. (1 point) What is the reducing agent in the following reaction? 2K + S - - > K2S *

• K

• S

• K2S

• None of the above

5. (1 point) The oxidation number of oxygen when it is in a compound other than a peroxide is ____. *

• -2

• -1

• 0

• +2

6. (1 point) What is the oxidation number for each atom in NH4F? *

• N=+3, H=-4, F=+1

• N=-3, H=+1, F=-1

• N=+3, H+1, F=-1

• N=-3, H=-1, F=+1

7. (1 point) Which atom has a change in oxidation number of +4 in the following redox reaction?K2Cr2O7 + H2O + S - - > KOH + Cr2O3 + SO2 *

• K

• Cr

• O

• S

Metal activity series for question 8

8. (1 point) Of the following metals, which ions are most easily reduced? Use the above table to answer the question. *

• magnesium

• lead

• copper

• sodium

9. (1 point) In a zinc-copper cell, Zn(s) | Zn^+2(1M) || Cu^+2(1M) | Cu(s), which electrode is positive? *

• Cu^+2

• Cu(s)

• Zn(s)

• Zn^+2

10. (1 point) How must a redox reaction that is to be used as a source of electrical energy be set up? *

• One half reaction must involve two metals.

• Two half reactions must use a metal wire electrodes.

• One half reactions must involve more than one electron.

• Two half reactions must be physically separated.

11. (1 point) In a voltaic cell at which electrode does reduction occur? *

• anode only

• cathode only

• both anode and cathode

• Either anode or cathode , depending on the metal

12. (1 point) What assures that there is no charge build up in a voltaic cell as oxidation and reduction occur? *

• one of the half cells

• the electrolyte solutions

• the moving electrons

• the salt bridge

13. (1 point) In the most common dry cell what material is the electrode that is the center rod made of? *

• copper

• graphite

• lead

• zinc

14. (1 point) To charge a lead storage battery you must ________. *

• apply a direct current to it.

• hold a magnet close to it.

• Use alternating current through it.

• increase the oxygen in it.

15. (1 point) In a hydrogen- oxygen fuel cell ______________. *

• hydrogen diffuses through the cathode.

• hydrogen and oxygen are mixed before entering the anode.

• oxygen diffuses through the anode.

• oxygen diffuses through the cathode.

16. (1 point) When electrical energy is used to cause a chemical reaction we call this process _____. *

• electronation

• electrolysis

• hydration

• reduction

17. (1 point) Which of the following is true about an electrolytic cell? *

• It changes chemical energy into electrical energy.

• It is the type of cell used in electrocution

• It uses an electric current to make a nonspontaneous reaction go.

• All of the above

18. (1 point) The products formed in the net reaction of the electrolysis of water are ____. *

• aqueous hydrogen ion and hydroxyl ion

• hydrogen gas and oxygen gas

• hydrogen gas and liquid oxygen

• water vapor

19. (1 point) What occurs in electroplating? *

• decomposition of an ionic compound

• decomposition of a salt layer

• deposition of a metal layer on a material

• deposition of a salt layer on a metal

Open response Questions : Answer questions 20 through 22 in complete sentences

20. ( 5 points) Write the equation for the reaction of calcium and oxygen ( 1 point). Write the oxidation numbers for both the reactants and the products ( 2 points) What is the reducing agent and why? ( 2 points)

21. ( 5 points) Tell how a voltaic cell works. Be sure to mention the various parts of the voltaic cell and what they have to do with how it works.

22. (4 points) What are the signs given to the electrodes in both the voltaic and electrolytic cells? What reaction (oxidation or reduction) is occurring at each?

To test for Ag+

1. add 2 drops 6M HCl to 1 mL of unknown sample. Centrifuge (3-4 minutes). Add 1 more drop of 6M HCl and centrifuge (if necessary). A precipitate will form; decant (pour off) solution to another test tube (save solution)

2. wash precipitate from step 1 with 1 or 2 mL of distilled water. Stir, centrifuge, and decant the liquid (can be discarded). Then add 1 mL of distilled water to precipitate and place the test tube in boiling water. Let the test tube boil, and stir occasionally. Centrifuge the hot test tube (solution) and then decant after centrifuging (save the decant solution). Save the precipitate.

3. Add 10 drops of 6M HNO3 to the precipitation from step 2 and stir it. Centrifuge and decant the solution (save it).

4. Add 6M HNO3 to the solution from step 3 until it is acidic toward litmus paper. If Ag+ is present in the acidic solution, a white precipitate of AgCl will form.

To test for Cu2+ and Sb3+

1. Add 15M NH3 to 1 mL of unknown sample until the solution is basic to litmus paper. A precipitate may form at this point due to the formation of insoluble oxychlorides of bismuth, tin, or antimony. Add 1 drop 6M HCL for each mL of solution. Verify the pH level by using methyl violet paper, the paper should turn blue-green. Adjustment may be necessary to achieve the blue-green color. Add HCl or NH3 until it is almost the same blue-green color.

2. Add 15-20 drops of 1M thioacetamide to solution and heat in boiling water for 5 minutes. Centrifuge and decant to another test tube (may contain Fe and Cr). Test solution for completeness of precipitation by adding 2 drops of thioacetamide and letting it sit for 1 minute. If precipitation forms, add a few more drops of thioacetamide solution and heat again. Combine the two batches of precipitate. Wash precipitate with 1 M NH4Cl solution, stir, and warm in bath water; centrifuge, and discard the wash solution.

3. To sulfide precipitate from step 2 add 15 drops of 6M KOH, 2mL of water, and 3 drops 1M thioacetamide. Heat solution in boiling water for 5 minutes and stir occasionally. Centrifuge and decant the solution. Wash any precipitate remaining twice with 5 drops of water and combine the washings with the solution. Tin and antimony will be present in the solution as complex ions, and copper and bismuth will remain as undissolved sulfide precipitate. Save precipitate for later.

4. Add 12 M HCL to the solution from step 3 until it is acidic toward litmus paper. When acidified, some tin and antimony will precipitate as orange sulfides. Add 5 drops of 1M thioacetamide and heat. Centrifuge and decant solution (this solution can be discarded)

5. Add 10-15 drops of 12M HCl to the precipitate from step 4 and heat for 5 minutes to dissolve precipitate. Continue heating until H2S is gone. Centrifuge out any insoluble sulfur residue.

6. Add 0.5 grams of oxalate acid and 2 mL of water to remainder of solution from step 5. Heat the solution, if necessary, to dissolve the oxalate acid. This reagent forms a stable complex ion with Sn4+. Add 10 drops of 1M thioacetamide and heat. Formation of an orange precipitate indicates the presence of antimony.

6. Add 30 drops of 6M HNO3 to remaining sulfide precipitate from step 3. Heat solution gently in boiling water to complete dissolution of the sulfides. Centrifuge the solution to remove sulfur and decant it into a small beaker. Place beaker on hot plate and evaporate the liquid to a volume of about one-half mL. DO NOT LET SAMPLE DRY OUT.

7. Add 15M NH3 dropwise to the solution obtained in step 8 until it is basic toward litmus paper. Add a few more drops. The formation of a deep blue solution indicates the presence of copper.

Testing for Fe3+ and Cr3+

1. [If a solution is being used from which group II cations have been precipitated, remove the H2S and excess acid by boiling the solution in a small beaker until a volume of about 1 mL. Remove any sulfur residue by centrifuging the solution.] [if a solution being used contains only group II ion, then use 1 mL of this solution.]

2. Add 3 drops of 3% H2O2 to 1 mL solution from step 1. Add 15 drops of 6M NaOH to solution to make it strongly alkaline (basic). Stir for 1 minute and boil to remove all excess H2O2. After H2O2 is decomposed, solution will ‘bumb’. If oxidation of chromium (II) has occurred, the solution will be yellow. If not, add 3 more drops of H2O2 and boil again. Centrifuge and decant. Save precipitate (contains Fe3+) and solution (contains Cr3+).

3. Acidify solution from step 2 by adding drops of 15 M HNO3(about 5 drops) and test with litmus paper. Add 6M NH3 drop by drop with mixing until solution is basic. This treatment will precipitate Al3+, will be gelatinous white. May appear yellow because of the presence of CrO42-. Centrifuge and decant the solution, it contains CrO42-. Add 1 M BaCl2 to solution from step 3 to precipitate yellow BaCrO4. If precipitate is slow to form, heat in boiling water. Centrifuge the solution and decant. Dissolve the precipitate in 3 drops of 6M HNO3; heat and stir. Add 4 drops of water and cool it down. Add 6 drops of diethyl ether followed by 1 drop of 3% H2O2. Shake gently, and allow layers to settle. A blue color in ether (top) later confirm test for chromium.

4. Add 8 drops of 15M HNO3 to precipitate from step 2. Heat until precipitate is dissolved. Allow to cool and add 10 drops of 1M NH4Cl and then add 15 M NH3 until solution is alkaline to litmus paper. This will precipitate iron as brown. Then add 4-5 drops of NH3 solution and stir. Centrifuge and decant solution (this solution may be discarded).

5. Dissolve the precipitate from step 5 by adding 10 drops of 6M HCl. Dilute with 2mL of distilled water and add 2 drops of 0.5M KSCN. A blood-red color indicates presence of Fe3+.

-----------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------

I will be given 3 mL of an unknown sample and I will have 75 minutes to do a qualitative analysis for Ag, Cu, Sb, Fe, Cr.

I need help coming up with a flow chart based on the above steps so that it is makes this class assignment smoother since I will only be given 75 minutes to complete it.

Thank you!

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