Help Completing this Ionic Equation

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The discussion centers on identifying whether the given equation is a redox reaction and how to complete it. Participants emphasize the importance of writing down the products to determine the reaction type. The equation involves three reactants: Cl^-, Ag(NH3)2^+, and H^+. To recognize a redox reaction, one must identify the oxidation and reduction half-reactions. Completing the equation is essential for further analysis and understanding of the reaction dynamics.
Jeff231
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Is this equation a redox reaction? If it is how do I recognize it? I can't seem to complete this.

If it is a redox equation, what are the two half reactions becuase there are three reactants?

Cl^{-}~+~Ag[(NH_{3})_{2}]^{+}~+~H^{+}==>~??

Thanks.
 
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It's not a reaction until you write down the products. What do you think will form?
 
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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