How to use standard reduction potential table

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Discussion Overview

The discussion revolves around understanding how to use a standard reduction potential table to identify oxidation and reduction half-reactions in redox reactions. Participants explore the application of these concepts in both voltaic and electrolytic cells, addressing specific examples and the conditions under which reactions occur.

Discussion Character

  • Exploratory
  • Technical explanation
  • Debate/contested
  • Mathematical reasoning

Main Points Raised

  • One participant questions how to identify oxidation and reduction half-reactions using the standard reduction potential table, citing a specific reaction involving calcium and bromine.
  • Another participant emphasizes the importance of balancing the reaction and suggests assigning oxidation numbers to clarify oxidation and reduction processes.
  • Some participants propose that the half-reaction with the higher standard potential will oxidize the one with the lower standard potential, providing examples with iron and copper.
  • One participant discusses the relationship between Gibbs free energy (ΔG°) and standard potentials, indicating that a negative ΔG° suggests a spontaneous reaction.
  • There is a discussion about the conditions under which zinc oxidizes copper, despite the standard potentials suggesting otherwise, raising questions about the context of the reactions.
  • Participants mention the necessity of having a positive cell potential for a voltaic cell and provide calculations based on standard reduction potentials.
  • One participant raises the scenario of non-voltaic cells and the implications of negative cell potentials, suggesting that the context of the problem is crucial for determining spontaneity.

Areas of Agreement / Disagreement

Participants express differing views on how to interpret the standard reduction potential table, particularly regarding specific examples. While some agree on the general principles of oxidation and reduction, the application to particular reactions remains contested and unresolved.

Contextual Notes

Participants note that the reactions discussed may not be balanced or may depend on specific conditions, such as whether the reactions are spontaneous or non-spontaneous. The discussion highlights the complexity of applying standard reduction potentials in various contexts.

Who May Find This Useful

This discussion may be useful for students and practitioners in chemistry, particularly those studying electrochemistry, redox reactions, and the application of standard reduction potential tables.

TheRedDevil18
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Hi, can someone please explain to me how I can identify which is the oxidation half and which is the reduction half of a redox reaction. I have read on the internet that the oxidation half is the highest up the table between the two then why doesn't it work for:

(Ca^2+) + 2Br → Ca + Br2, calcium is higher up the table but it is being reduced. Using table 4B.
Am I using the table wrong?
 
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Where did you got the reaction equation from?

Apart from everything else, it is not balanced.
 
I don't see anything wrong with the balancing, 1 calcium and two bromines on both sides
 
What about charge?
 
Sorry, (Ca^2+) + (2Br^-) → Ca + Br2
Now could you please answer my question
 
Such reaction would not happen.

Other than that - assign oxidation numbers to all atoms. Oxidation means ON goes up, reduction means ON goes down.
 
I know that but how would I know which is oxidation and which is reduction by using a half reaction table
 
I guess now we are getting to the point.

Generally speaking the one with the higher standard potential will oxidize the one with the lower standard potential.

So for example, we have two half reactions:

Fe3+ + e- <-> Fe2+ (E0 = +0.77V)

Cu2+ + 2e- <-> Cu(s) (E0 = +0.34V)

Fe3+ will oxidize metallic copper to Cu2+.
 
I just always work it out by keeping in mind that you need ΔG° < 0 for a reaction to proceed spontaneously and that ΔG° = -nFE°. So when you look at a table of standard reduction potentials you need to flip the half reactions so as to get the most positive E° (when added together) which will then give the most negative ΔG°, showing that the reaction proceeds as written (after balancing etc).
 
  • #10
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Borek said:
I guess now we are getting to the point.

Generally speaking the one with the higher standard potential will oxidize the one with the lower standard potential.

So for example, we have two half reactions:

Fe3+ + e- <-> Fe2+ (E0 = +0.77V)

Cu2+ + 2e- <-> Cu(s) (E0 = +0.34V)

Fe3+ will oxidize metallic copper to Cu2+.

Okay, then what about:

(Zn^2+) + 2e → Zn (-0.76)
(Cu^2+) + 2e → Cu (+0.34)

By standard potential you mean the voltages?, because Cu voltage is higher than Zn but yet Zn oxidises Cu
 
  • #11
That table is handy, you can quickly see which substances can be combined into working voltaic cells. For you, it shows the calcium related half reaction is -2.87 V. The bromine related one is +1.07 V. For this to be a working voltaic cell, the total cell potential must be positive (=main idea for these types of problems)

E° (cell) = E°(cathode) - E°(anode) (> 0 for voltaic)

With those 2 numbers from the table the only way that can happen (only way to make a positive difference) is if bromine is the cathode and calcium the anode.

+1.07 - (-2.87) = +3.94 V > 0

Therefore bromine is getting reduced and calcium getting oxidized (Reduction and cathode both start with consonants, oxidation and anode both start with vowels, is how I remember).

So your equation would actually be

Br2 + Ca (s) -> Ca2+ (aq) + 2Br-(aq)

This shows bromine gaining electrons and calcium losing them.

Now for your new example with the Zn (-0.76) and Cu (+0.34), the only way to get a working voltaic cell is if Zn is the anode and copper the cathode. Otherwise you would get a negative number. Then again for when writing the equation, you just need to show Zn being oxidized (turning into cations) and copper being reduced (going from ions into a solid).

Hope that helps.
 
Last edited:
  • #12
TheRedDevil18 said:
Okay, then what about:

(Zn^2+) + 2e → Zn (-0.76)
(Cu^2+) + 2e → Cu (+0.34)

By standard potential you mean the voltages?, because Cu voltage is higher than Zn but yet Zn oxidises Cu

In standard conditions?
 
  • #13
mishima said:
That table is handy, you can quickly see which substances can be combined into working voltaic cells. For you, it shows the calcium related half reaction is -2.87 V. The bromine related one is +1.07 V. For this to be a working voltaic cell, the total cell potential must be positive (=main idea for these types of problems)

E° (cell) = E°(cathode) - E°(anode) (> 0 for voltaic)

With those 2 numbers from the table the only way that can happen (only way to make a positive difference) is if bromine is the cathode and calcium the anode.

+1.07 - (-2.87) = +3.94 V > 0

Therefore bromine is getting reduced and calcium getting oxidized (Reduction and cathode both start with consonants, oxidation and anode both start with vowels, is how I remember).

So your equation would actually be

Br2 + Ca (s) -> Ca2+ (aq) + 2Br-(aq)

This shows bromine gaining electrons and calcium losing them.

Now for your new example with the Zn (-0.76) and Cu (+0.34), the only way to get a working voltaic cell is if Zn is the anode and copper the cathode. Otherwise you would get a negative number. Then again for when writing the equation, you just need to show Zn being oxidized (turning into cations) and copper being reduced (going from ions into a solid).

Hope that helps.

What if it where non-voltaic or assuming I don't know whether it is voltaic or non-voltaic?

Borek said:
In standard conditions?

Yes, I think so
 
  • #14
If it were an electrolytic cell you would be trying to satisfy the condition that E (cell) < 0. This implies the reaction is non-spontaneous.

Usually you know what kind of cell you are dealing with. If not, the textbook will give you clues, such as providing the redox reaction. The standard kind of textbook problem is something like "determine if the following is a spontaneous or non-spontaneous reaction", that is the same as asking "determine if the following is a voltaic or electrolytic cell" (at least in the context of electrochemistry), and is the same as "determine if the cell potential is positive or negative". For that type of problem, you would be given half-reactions or commonly a full redox equation. If you are using a text, take a look at the end of chapter problems and see if you have questions with that pattern. In real life, you would probably already know since you could just look for a voltage source, or even probe with a voltmeter.

For both types: Oxidation is always anode, reduction is always cathode. Electrons always flow from anode to cathode.
 

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