Balancing Redox Reactions (another question) help?

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To balance the redox reaction Zn + HNO3 → Zn(NO3)2 + NH4NO3 + H2O, it is essential to break it down into half reactions. Start by determining the oxidation states of all elements to identify which substances are oxidized and reduced. The oxidation number increases for the oxidized species and decreases for the reduced species. While the oxidation number method can be complex, it is a reliable way to find the necessary changes for balancing. Understanding these concepts will facilitate the correct application of half reactions in the balancing process.
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Homework Statement


Use half reactions to balance the following redox reactions and underline the oxidizing agent
m) Zn + HNO3 --------> Zn(NO3)2 + NH4NO3 + H2O

Homework Equations


Not applicable

The Attempt at a Solution


I would usually start by breaking down the equation into half reactions, but I'm not actually sure how to do this with this equation, since I have an H2O in there.
 
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If you are not sure what is getting oxidized and what is getting reduced, start assigning oxidation numbers to all elements on both sides of the reaction equation.
 
Borek said:
If you are not sure what is getting oxidized and what is getting reduced, start assigning oxidation numbers to all elements on both sides of the reaction equation.

Yeah, but the question asks me to use half reactions to balance the redox equation, not oxidation numbers. Is there a way to do this using half reactions?
 
There is no problem with balancing using half reactions, but first you have to find out what is being reduced and what is being oxidized. If you don't see it at the first sight, ON are a sure way of finding out.
 
Borek said:
There is no problem with balancing using half reactions, but first you have to find out what is being reduced and what is being oxidized. If you don't see it at the first sight, ON are a sure way of finding out.

How would you use the oxidation numbers to determine what is being oxidized and reduced? Is it to do with the change, so if the oxidation number increases you have an oxidation and if it decreases you have a reduction? I haven't really learned the oxidation method for balancing redox equations in school yet, so I'm still kind of iffy about doing that. Could you explain to me how to determine this?
 
Thread 'Confusion regarding a chemical kinetics problem'

Thread 'Confusion regarding a chemical kinetics problem'

TL;DR Summary: cannot find out error in solution proposed. [![question with rate laws][1]][1] Now the rate law for the reaction (i.e reaction rate) can be written as: $$ R= k[N_2O_5] $$ my main question is, WHAT is this reaction equal to? what I mean here is, whether $$k[N_2O_5]= -d[N_2O_5]/dt$$ or is it $$k[N_2O_5]= -1/2 \frac{d}{dt} [N_2O_5] $$ ? The latter seems to be more apt, as the reaction rate must be -1/2 (disappearance rate of N2O5), which adheres to the stoichiometry of the...
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