I am stuck on #29.  First of all, I am not sure with the question about how many g of water was required, I am supposed to use the 45g of NO2 or the NO2 that actually reacted (39.03g).  Same for theoretical yield.  Do I use the value that should have reacted, or I expect based on what is left over?  Can I assume water was not limiting?  Even with all my possible answers, I can't get the value on the answer key.  I am missing something.  Can someone help me?

Hello,

Not sure if I am looking at the correct problem since I don't see 45 g of NO2 anywhere and I don't see water involved in the reaction.

See if my explanation below helps:

For part b of #29:  convert the 0.0065 g NO2 to moles. Use the resulting number to find each, the moles of oxygen needed and the moles of NO needed.

For part c of #29:  the 3 moles of NO2 represent how much formed at 86% yield. The first step is to find out how much would have been produced at 100% yield. (100% yield = 3/0.89 = 3.37 moles NO2 produced). Use that value to determine the moles of oxygen needed. Lastly, use the relationship 22.4 L = 1 mole of gas at 1 atm and 273K to convert moles to liters.

Please be sure to let us know if this does not help you with this problem or if you have any further questions.

This is the question in the worksheet. I am sorry.  I think its supposed to be number 28.  I must have renumbered it on my worksheet for my students.  
In excess oxygen, 45 g of nitrogen dioxide reacted with water and 5.97 g of nitrogen 
 	  dioxide was left over.
4 NO2(g) + O2(g) + 2 H2O(l) → 4 HNO3(aq)
How many grams of water was required to react?
So based on left over NO2, I deduce that 39.03g of NO2 reacted.  Therefore, the amount of water that reacted should be 7.63g H2O (this is same as the answer key)
What would be the theoretical yield?
Should the theoretical yield be based on how much NO2 was used or how much NO2 that reacted?  Are we assuming that water is limiting or NO2 is limiting and the amount of NO2 reacted is due to inefficiency? 
Based on 45g NO2 the theoretical yield should be 61.6g HNO3
Based on 35.03g NO2 the theoretical yield should be 53.5g HNO3  The answer key is twice this. 
If 80 grams of nitric acid was produced, what was the percent yield for the reaction?
I can not calculate this because based on my calculations, I didn’t get 80g.

Also, logically, I am thinking its a lot to get 80g product from 45g of reactant when product and the reactant has 1:1 ratio.