For the following electrochemical cell, what is the correct, balanced reaction?

Cu(s) | Cu2+(aq) || Ag+(aq) | Ag(s)

A. Cu(s) + Cu2+(aq) → Ag+(aq) + Ag(s)
B. Cu(s) + Ag+(aq) → Cu2+(aq) + Ag(s)
C. Cu(s) + 2 Ag+(aq) → Cu2+(aq) + 2 Ag(s)
D. Cu2+(aq) + Ag(s) → Cu(s) + Ag+(aq)

My guess is Answer D

For the following electrochemical cell, what is ΔEº?
Zn(s) | Zn2+(aq) || Cr3+(aq), Cr2+(aq) |Pt(s)

Zn2+(aq) + 2 e- → Zn(s) Eº = -0.8 V
Cr3+ (aq) + e- → Cr2+(aq) Eº = -0.4 V

A. 0.4 V
B. 0.0 V
C. -0.4 V
D. -1.2 V

I am guess it is C..but I am pretty confused about how to solve this--if someone could explain this that'd be great. I changed the sign of the 2nd half reaction so it would be an oxidation (i think..) then added...but I could really use clarification.

Which of the following reagents is most likely to function as a reducing agent?

A. F2(g) F2(g) + 2 e- → 2 F-(aq) Eº = 2.9 V
B. H2(g) 2 H+(aq) + 2 e- → H2(g) Eº = 0 V
C. Al(s) Al3+ (aq) + 3 e- → Al(s) Eº = -1.7 V
D. They are all equally likely

I am guessing it would be C because it has more negative standard reduction potential, so there will be oxidation and it will be a redcing agent? I keep getting mixed up...Thanks for any help

Your "guess" on 1 is correct, but why guess? I suppose you were just kidding and that you worked it out.

On #2, no. You wrote the correct equation for the first one, now do that for the second one.
Zn + 2Cr^+3 ==> Zn^+2 + 2Cr^+2
Zn ==> Zn^+2e Eo as an oxidation is +0.8
Cr^+3 + e ==> Cr^+2 Eo as a redn is -0.4
Add them together.
Ecell = Eo(oxdn) + Eo(redn) = +0.4
A + Eo means the reaction takes place spontaneously. Eo is - means the cell will not go spontaneously but will if the reactions are reversed.

On the third one you are correct. It IS easy to get confused and I do, too, because I learned all of this with all of my tables written as oxidations and not reductions. So I cheat a little. Since my tables run from a high + voltage to a more negative - voltage, the ones at the top Zn, Al, etc, are reducing agents. So I turn the three reactions around in my mind and look for the one with the largest + sign. That is Al. You look for the one with the most negative sign. But you're doing it right.

Which of the following are correct symbols and charges? Could be as many as 5.

1. Na-
2. Ag+
3. O2-
4. Ag-
5. NO3-
6. Cu2+
7. F+
8. Cl2-

For the first question:

To determine the correct, balanced reaction in an electrochemical cell, you need to identify the half-reactions for the oxidation and reduction occurring at each electrode. In this case, the copper electrode is being oxidized, while the silver electrode is being reduced.

The half-reactions are as follows:

Oxidation (anode): Cu(s) → Cu2+(aq) + 2e-
Reduction (cathode): Ag+(aq) + e- → Ag(s)

When combining these two half-reactions, you need to ensure that the number of electrons transferred is the same in both reactions. Multiply the second half-reaction by 2 to balance the electrons:

2Ag+(aq) + 2e- → 2Ag(s)

Now you can combine the two half-reactions to form the balanced equation, taking into account the number of electrons transferred:

Cu(s) + 2Ag+(aq) → Cu2+(aq) + 2Ag(s)

Therefore, the correct balanced reaction is Option C: Cu(s) + 2 Ag+(aq) → Cu2+(aq) + 2 Ag(s).

For the second question:

To determine ΔEº (the standard cell potential), you need to subtract the standard reduction potential of the anode from the standard reduction potential of the cathode.

Given the following information:
Zn2+(aq) + 2 e- → Zn(s) Eº = -0.8 V (anode)
Cr3+ (aq) + e- → Cr2+(aq) Eº = -0.4 V (cathode)

Since you're subtracting the anode potential from the cathode potential, the sign of the anode half-reaction potential should be flipped (as you mentioned).

Now subtract the anode potential from the cathode potential:

ΔEº = Eº(cathode) - Eº(anode)
= (-0.4 V) - (-0.8 V)
= -0.4 V + 0.8 V
= 0.4 V

Therefore, the correct answer is Option A: 0.4 V.

For the third question:

To determine which reagent functions as the reducing agent, you need to compare the standard reduction potentials (Eº).

The more negative the standard reduction potential, the more likely a species is to undergo reduction (be a reducing agent).

Comparing the reduction potentials given:

F2(g) Eº = 2.9 V
H2(g) Eº = 0 V
Al(s) Eº = -1.7 V

Since Al(s) has the most negative standard reduction potential, it is the most likely to function as a reducing agent. This is because it has a stronger tendency to undergo oxidation and lose electrons during a redox reaction.

Therefore, the correct answer is Option C: Al(s).