I don't know that there is a chemistry reason for it. First, I think its a matter of not wasting NaHCO3 although that isn't very expensive. Also, 5% NaHCO3 is a "kinda" lab standard for removing acids in the organic lab. Third is that using saturated NaHCO3 solution runs the risk of having some of the solid ppt when you don't want it to ppt (I know it is being used up so it won't be saturated for long BUT why take chances with a huge excess like that?) and finally, someone may have wanted to give you some practice at making up a 5% solution. Again, there may be a chemical reason but I don't know it if there is.
Is it due to it producing water in the rxn product?
I wouldn't think so. The amount of water formed depends upon the amount of excess CH3COOH, not on the amount of NaHCO3. So the same amount of water will be formed if you used 1%, 5%, or 100% NaHCO3.
But if I have to use excess acetic acid wouldn't that produce more water than another rxn which does not?
I have to take out the water with anhydrous sodium sulfate after extracting with the water & sodium bicarbonate solution....but I guess that woudn't make a difference since the water produced would just form a single layer with the water added for extraction.
I'm puzzled as to why they use 5% sodium bicarbonate.
Thanks for your help Dr.Bob
That's right. The water from the NaHCO3 solution (of whatever concentration) doesn't know the difference from the water in the acetic acid (you didn't add glacial I'm sure) and neither of those know the difference in the water from the neutralization reaction of the bicarbonate and the acid. You simply take care of the water with a drying agent. My best guess is,
I think it's just handy.
I'm not sure I want to state that 5% sodium bicarbonate is "handy", Dr.Bob.(this question is on my prelab assignment)
I seriosly think my lab professor would laugh. ^____^
Oh and Thank you for your help Dr.Bob =D