Partial pressure question (should be easy)

AI Thread Summary
The discussion revolves around calculating the partial pressure of products in a chemical reaction, specifically H2 + 2O2 = 2HNO2. Given the partial pressures of H2 at 0.5 and O2 at 0.3, it is clarified that the partial pressure of 2HNO2 cannot simply be summed to 0.8 due to the stoichiometric relationships in the equation. The importance of maintaining the correct mole ratio and recognizing that nitrogen is present only on one side of the equation is emphasized. Participants note that without equilibrium considerations, the calculation must adhere to the principles of stoichiometry. Accurate understanding of partial pressures in chemical reactions is crucial for correct assessments.
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Homework Statement



if you have an equation such as H2+2O2=2HNO2 and you know the partial pressure of H2 is .5 and the partial pressure of O2 is .3, does that mean that the partial pressure of 2HNO2 is .8?

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The Attempt at a Solution

 
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Assuming constant temperature and volume ; in the case there is no need to account for equilibrium you need to find the correct mole ratio by observing the stoichiometric equation.
 
As it was already signalled you can't balance equation that have nitrogen on only one side.
 
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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