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11 Dec 2019

Use a Standard Reduction table to write a balanced redox equation, and to determine the cell potential and the number of electrons transferred for the spontaneous process that occurs. Use no special characters for subscripts and superscripts and include the coefficient with the substance. 3Fe2+ should be written as 3Fe2+. Note that Ox1 and Red2 are the oxidizing and reducing agents, and Red1 and Ox 2 are the other half of their couples.

a) An Fe2+/Fe couple is connected to Sn2+/Sn couple.

Ox1 + Red2 Red1 + Ox2
1 [Correct: Your answer is correct.] + 2 [Correct: Your answer is correct.] 3 [Correct: Your answer is correct.] + 4 [Correct: Your answer is correct.]

The standard cell potential is E =

5 [Correct: Your answer is correct.] V

The number of electrons transferred is n =

6 [Correct: Your answer is correct.]

b) An Ag1+/Ag couple is connected to Sn2+/Sn couple.

Ox1 + Red2 Red1 + Ox2
7 [Correct: Your answer is correct.] + 8 [Correct: Your answer is correct.] 9 [Correct: Your answer is correct.] + 10 [Correct: Your answer is correct.]

The standard cell potential is E =

11 [Correct: Your answer is correct.] V

The number of electrons transferred is n =

12 [Correct: Your answer is correct.]

c) Aluminum is placed in hydrochloric acid.

Ox1 + Red2 Red1 + Ox2
13 + 14 15 + 16

The standard cell potential is E =

17 V

The number of electrons transferred is n =

18

d) Copper is placed nitric acid.

Ox1 + Red2 + Other Red1 + Ox2 + Other
19 + 20 + 21 22 + 23 + 24

The standard cell potential is E =

25 V

The number of electrons transferred is n =

26

Please help I was able to figure out the first two, but the last two are a bit challenging. Thank you in advance! P.S. the numbers that go up to 26 aren't answers they are the question numbers from webassign normally hidden, I tried to remove them to make it neat, but I was not allowed to post my question without them, sorry for the lack of organization.

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Related questions

Write net equation for the spontaneous redox reaction and determine the standard cell potential that results when each of the following redox couples are connected in an electrochemical cell. Instructions:

Enter all substances in the order listed at the top of the column.

Use a carot to indicate a superscript, but do nothing for subscripts.

Use a hyphen + greater than (->) for yields.

Click on the eye symbol to check your formatting.

Couples Rxn: Ox(1) + Red(2) + other -> Red(1) + Ox(2) + other Eo (V)
Cu/Cu2+ and Ag1+/Ag 2Ag^1+ + Cu -> 2Ag + Cu^2+ CORRECT 0.46
Pb2+/Pb and Fe3+/ Fe2+ 2Fe^3+ + Pb -> 2Fe^2+ + Pb WRONG-HELP 0.9
H1+/H2 and Ni2+/Ni 2H+ + Ni -> H^2 + Ni^2+ WRONG HELP 0.23
Cr/Cr3+ and Ti2+/Ti 2Cr^3+ + 3Ti -> 2Cr + 3Ti^2+ CORRECT 0.89

The two middle reactions are wrong. Not sure if the formulas are wrong or if I am just entering it incorrectly?

± Cell Potential and Equilibrium
18 of 28
Constants | Periodic Table
The equilibrium constant, K, for a redox reaction is related to the standard potential, E?, by the equation
lnK=nFE?RT
where n is the number of moles of electrons transferred, F (the Faraday constant) is equal to 96,500 C/(mol e?) , R (the gas constant) is equal to 8.314 J/(mol?K) , and T is the Kelvin temperature.
Standard reduction potentials
Reduction half-reaction E? (V)
Ag+(aq)+e??Ag(s) 0.80
Cu2+(aq)+2e??Cu(s) 0.34
Sn4+(aq)+4e??Sn(s) 0.15
2H+(aq)+2e??H2(g) 0
Ni2+(aq)+2e??Ni(s) ?0.26
Fe2+(aq)+2e??Fe(s) ?0.45
Zn2+(aq)+2e??Zn(s) ?0.76
Al3+(aq)+3e??Al(s) ?1.66
Mg2+(aq)+2e??Mg(s) ?2.37

Part A Part complete
Use the table of standard reduction potentials given above to calculate the equilibrium constant at standard temperature (25 ?C) for the following reaction:
Fe(s)+Ni2+(aq)?Fe2+(aq)+Ni(s)
Express your answer numerically.
View Available Hint(s)
K = 2.68×10^6 Correct
When E?>0 and K>1 the reaction favors the products.

Part B
Calculate the standard cell potential (E?) for the reaction
X(s)+Y+(aq)?X+(aq)+Y(s)
if K = 7.87×10?3.

Express your answer to three significant figures and include the appropriate units.

E? = ?

experment1 Electron Transfer (Redox) Reactions

1. Comparison of Br2/Br– and Fe3+/Fe2+ Couples

Oxidant-reductant pairs from these couples will he tested, i.e. Br2 with the reductant Fe2+ and

Br– with the oxidant Fe3+. A reaction will occur in only one of these cases because of the

difference in electrode potential between the two couples.

To test whether reaction has occurred, SCN– will be used as a probe for Fe3+, with which it

gives a red complex. Note that the air rapidly oxidises Fe2+, so that a faint pink colour is often

obtained when SCN– is added to Fe2+.

Solutions containing Fe2+ or Fe3+ tend to decompose (although in different ways); both are

stabilised by H+. The iron(II) solution provided is freshly prepared 0.1 M (NH4)2Fe(SO4)2 / 2

M H2SO4.

To 5 drops of the 0.1 M iron(II) solution in a semi-micro tube add 10 drops of 0.1 M KBr and

stir. Place 1 drop of the mixture on a white tile, and test for Fe3+ by adding 1 drop 0.1 M

KSCN. A strong red colour indicates Fe3+.

Repeat using 5 drops of the 0.1 M iron(III) solution and adding 10 drops of 0.1 M KBr.

Repeat using 5 drops of the 0.1 M iron(II) solution and adding 10 drops of 0.1 M Br2.

Repeat using 5 drops of the 0.1 M iron(III) solution and adding 10 drops of 0.1 M Br2.

2.Use of the Br2/Br– and Fe3+/Fe2+ Couples in an Electrochemical Cell

You are supplied with a pair of carbon rods joined by copper wire. Place two 100 mL beakers

side by side.

In one put 50 mL of the iron(II) solution and 1 mL of 0.1 M KSCN solution. In the other

beaker put 25 mL of 0.2 M Br2 solution and 25 mL of 4 M HCl.

Correct any difference in level between the solutions in the two beakers by adding water to

one. Fold an 11 cm filter paper into a strip about 1 cm wide. Bend the strip in the middle,

moisten with a little 4 M HCl, then insert the strip so that it links the two solutions. Insert one

carbon rod into each solution, then allow to stand undisturbed for 30 minutes while

proceeding with Experiment 3. Carefully observe the apparatus at intervals and record any

change that occurs.

3. Comparison of the Fe3+/Fe2+ and Sn4+/Sn2+ Couples

In this experiment the additional sources of ions are:

tin(IV) ion 0.1 M SnCl4/1 M HCl

tin(II) ion 0.1 M SnCl2/1 M HCl

(Both Sn2+ and Sn4+ solutions are stabilised by H+.)

To 5 drops of the 0.1 M iron(II) solution in a semi-micro tube add 10 drops of the appropriate

0.1 M tin solution and stir. Place 1 drop of the mixture on a white tile, and test for Fe3+.

Record the results.

Similarly, test the 0.1 M iron(III) solution with an excess of the appropriate 0.1 M tin

solution.

You have now enough information to state whether the standard electrode potential of the

Fe3+/Fe2+ couple is more positive or less positive than that of the Sn4+/Sn2+ couple.

please answer following questions

1.Comparison of Br2/Br– and Fe3+/Fe2+ couples

Complete the following table.

Solutions

Observation

Fe3+ present (Yes/No)

Did a reaction occur?

Fe(II) and KBr

Fe(III) and KBr

Fe(II) and Br2

Fe(III) and Br2

2.Br2/Br– and Fe3+/Fe2+ couples in an Electrochemical Cell

a)What observations did you make over 3 minutes? Did Br2 or Fe2+ react? Write the ion/electron half- equations for this reaction and the complementary reduction. Write the overall redox reaction.

b)Draw/scan/insert a circuit diagram showing the half-reactions that occur in each beaker, the direction of electron flow and the direction of flow of ions.

3.Fe3+/Fe2+ and Sn4+/Sn2+ couples

Complete the following table.

Solutions

Observation

Fe3+ present (Yes/No)

Did a reaction occur?

Fe(II) and Sn(IV)

Fe(II) and Sn(II)

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