How Do You Balance Redox Reaction Equations?

Click For Summary

Discussion Overview

The discussion revolves around the process of balancing redox reaction equations, specifically focusing on the oxidation and reduction of silver and nitrate in an acidic solution. Participants explore the concepts of half-reactions, the roles of oxidizing and reducing agents, and the steps necessary to achieve a balanced equation.

Discussion Character

  • Homework-related
  • Technical explanation
  • Debate/contested

Main Points Raised

  • Some participants express confusion about balancing redox reactions and seek clarification on the differences between oxidizing and reducing agents.
  • There is a discussion about the half-reaction method, with some participants stating that half-reactions must show either reduction or oxidation.
  • Participants propose that silver is oxidized to Ag+ and that nitrate acts as the oxidizing agent, but there is uncertainty about the correct half-reaction for nitrate.
  • One participant suggests that the oxidation state of silver increases from 0 to +1, confirming its role as a reducing agent.
  • There are conflicting views on the correct half-reaction for nitrate, with some participants suggesting it reduces to NO while others express doubt about this outcome.
  • Several participants emphasize the importance of balancing charges and atoms in the half-reactions and the overall equation.
  • One participant provides a detailed breakdown of the half-reactions and the net balanced equation, but others question the accuracy of the proposed steps.

Areas of Agreement / Disagreement

Participants generally agree on the basic concepts of oxidation and reduction but express differing views on the specific half-reactions and the balancing process. The discussion remains unresolved regarding the correct formulation of the nitrate half-reaction and the overall balanced equation.

Contextual Notes

Participants highlight the need to account for charges and the correct oxidation states when balancing the equations. There is mention of the reaction occurring in an acidic solution, which may affect the balancing process.

Who May Find This Useful

This discussion may be useful for students learning about redox reactions, particularly those struggling with the concepts of oxidation, reduction, and the half-reaction method in chemistry.

southerngirl5390
Messages
14
Reaction score
0
Ok I'm trying to do my chemistry homework and well I'm stuck. I have the first few questions done I'm stuck on question 5 step 3, an I'm not sure the other are correct. I still have no idea how to balance equations and well here we go again with balanceing things...

1. What is the difference between an oxidizing agent and a reducing agent?
Reduction means what it says: the oxidation number is reduced in reduction. Oxidation is the reverse process: the oxidation number of an atom is increased during oxidation.
2. When first learning to balance equations, we learned that the number of atoms of each element in the products and reactants must be equivalent. What are some additional factors that must be taken into account when balancing equations for redox reactions?

Combine the half reactions to eliminate the electrons from the overall reaction.

3. What are half reactions?
A half-reaction is simply one which shows either reduction OR oxidation, but not both.

4. What two aspects of the half-reaction equations must be balanced?
atoms and charge in order

5. For the equation Ag + NO3 - à Ag + + NO
(Note: This reaction takes place in an acidic solution.)

Step 1: What substance is reduced?
Oxygen
Step 2: What substance is oxidized?
Silver
Step 3: What is the half reaction for oxidation?

Step 4: What is the half reaction for reduction?

Step 5: What is the net balanced equation?

Step 6: What is the reduced equation?
 
Physics news on Phys.org
This looks like silver metal is being oxidized to Ag+. The oxidizing species is nitrate (acidic). Can you write a half reaction for what happens to silver?
 
ag-->ag+ ?? I'm not sure if that's what you mean but what happens to silver is it is oxidized ag-->ag+...? correct?
 
Where did the electron go? Show everything. Example

Fe -----> Fe+2 + 2e- (Half reaction)

You must account for the nitrate half reaction as well. A good place to start is to write down an equation that is balanced.
 
You mean like this?
Ag---->Ag+ 2g-
and
NO3---->NO+O3
 
Red. agents are always oxidised
Oxi. agents are always reduced...right

since metals are e- donors...they are reducing agents! here Ag is the reducing agent. oxidation state(OS) of Ag is 0...that of Ag+ is +1... this confirms oxidation of Ag, i.e oxidation state has increased.

NO3- acts as an oxidising agent... OS of N in NO3- is +5 and that in NO is +2
there has been a decrease in oxidation number... this confirms reduction of NO3- to NO, i.e. OS of Nitrogen has decreased

oxidation... Ag -------> Ag+ + e-

reduction NO3- + e- ------> NO

wait a sec... are you sure of the above eqn? the half eqn for oxidising eqn is ok...but for the reducing one... can't find out...sry


the NO3- is supposed to accept e-,,, but NO cannot be formed... well I'm not sure
 
No I'm not sure of anything at this point all i know is what my teacher told me. Thats why I'm hopeing someone can help me and my chemically challenged self...
For the equation Ag + NO3 - à Ag + + NO
(Note: This reaction takes place in an acidic solution.)

Step 1: What substance is reduced?

Step 2: What substance is oxidized?

Step 3: What is the half reaction for oxidation?

Step 4: What is the half reaction for reduction?

Step 5: What is the net balanced equation?

Step 6: What is the reduced equation?
 
southerngirl5390 said:
You mean like this?
Ag---->Ag+ 2g-
and
NO3---->NO+O3

Did you just misspell 'e-'?

Balance the nitrate reaction so that there are as many oxygens on the left as the right. Show charges! NO3 is not correct. It's NO3-. You MUST account for all of these charges to even begin to answer this question correctly.

Check your work carefully!
 
Last edited:
Ag---->Ag+2e-
and
NO3- ----> NO+3e-

?
 
  • #10
refer to post #8.
 
  • #11
NO3- has nitrogen in the +5 oxidation state. NO has the nitrogen in the +2 state. This is a 3 electron reduction of nitrogen. The electrons come from the silver metal. Silver is oxidized from a zero valency to a +1.

Ag --------> Ag+ + 1e- (Ag oxidized, half reaction)

NO3- + 3e- ---------> NO + 2O-2

It was done in acidic solution and those two O-2's are just dying to be water molecules. Add some H+ to the nitrate half reaction to react with the O-2's.

NO3- + 3e- + 4H+ --------> NO + 2H2O (NO3- reduced, half reaction)

Multiply the silver half reaction by three to account for the 3e-'s in the above reaction and you get:

3Ag --------> 3Ag+ + 3e-

combining the two reactions yields:

3Ag + NO3- + 3e- + 4H+ ---------> 3Ag+ + 3e- + NO + 2H2O (net balanced equation)

Canceling the 3e-'s from both sides of the reaction yields:

3Ag + NO3- + 4H+ -----------> 3Ag+ + NO + 2H2O (reduced equation)
 

Similar threads

Chemistry How do we interpret the steps in the algorithm to balance a redox equation?
  • · Replies 7 ·
Replies
7
Views
3K
Chemistry "The oxidized species of the couple is very oxidizing". Does this make sense?
  • · Replies 5 ·
Replies
5
Views
2K
Chemistry What is the purpose of redox reactions?
  • · Replies 18 ·
Replies
18
Views
5K
Chemistry How to apply Nernst equation in this problem?
  • · Replies 3 ·
Replies
3
Views
3K
Chemistry How to find the pH of a galvanic cell (MIT OCW problem set)
  • · Replies 2 ·
Replies
2
Views
4K
Can Methane and Citric Acid be Balanced in a Redox Reaction?
  • · Replies 3 ·
Replies
3
Views
2K
Is 3FeCl2 to 2FeCl3 and Fe(s) a Disproportionation Redox Reaction?
  • · Replies 1 ·
Replies
1
Views
3K
Number of electrons to balance redox reaction
  • · Replies 5 ·
Replies
5
Views
3K
Balance Redox Reaction: 2PbO2 + 4H+ = 2Pb 2+ + O2 + 2H2O
  • · Replies 1 ·
Replies
1
Views
3K
Redox reaction-naming the substance reduced
  • · Replies 4 ·
Replies
4
Views
2K