[yyttr3]

I was wondering if my calculations were correct, I started out with 3 cups of Vinegar at 5% acidity, I assume that means 5% acetic acid. Which means there should be .05*3=.15 cups of acetic acid in the vinegar or 35.488cm^3.

I believe acetic acid has a density of 1.05(g/cm^3), so there should be 37.26 gams of acetic acid in 3 cups of vinegar. With a molar mass of 60.05(g/mol) there should be 1.6116 mols of acetic acid.

If I have the formula right, then the reaction between sodium bicarb and acetic acid should be 1:1 and I need 1.6116 mols of sodium bicard. With a density of 2.20(g/cm^3) and  a molar mass of 84.07(g/mol) I should need .26cups of sodium bicarb.

Is these calculations correct? I'm trying to make sodium acetate for another experiment and I don't know if unreacted sodium bicarb will effect the experiment.

[Arkcon]

I was wondering if my calculations were correct, I started out with 3 cups of Vinegar at 5% acidity, I assume that means 5% acetic acid. Which means there should be .05*3=.15 cups of acetic acid in the vinegar or 35.488cm^3.

I'd prefer if you used SI units from the very beginning, rather than converting from cups to cm³ later, as you appear to.  At any rate, that is the conversion result I got from Goggle conversion.  So you're OK.  Warning -- you are ignoring significant figures, presumably because there is no instructor to rap your knuckles, and that is not appropriate.  You have 3 cups -- how close to the line were you?  You have 5% acetic acid -- what does that mean?  It may be perfectly fair for the company to sell 4.6% vinegar as 5% -- you really don't know.  So maybe only 35 mls of acetic can be said to be there with any confidence.

I believe acetic acid has a density of 1.05(g/cm^3), so there should be 37.26 gams of acetic acid in 3 cups of vinegar. With a molar mass of 60.05(g/mol) there should be 1.6116 mols of acetic acid.

Cool.  Here's a good time to ask, this was white vinegar, right?  But yeah.  Good math.

If I have the formula right, then the reaction between sodium bicarb and acetic acid should be 1:1 and I need 1.6116 mols of sodium bicard. With a density of 2.20(g/cm^3) and  a molar mass of 84.07(g/mol) I should need .26cups of sodium bicarb.

If the density is correct, I suppose you can do it that way.  It would be better to mass out the sodium bicarbonate on a gram scale.  I assume this was USP grade sodium bicarbonate.  That means it won't hurt a person who eats it, as it contains no toxic impurities as defined in the US Pharmacopeia.  That doesn't mean that it is chemically pure, it may be 99% pure, or even less.

Is these calculations correct? I'm trying to make sodium acetate for another experiment and I don't know if unreacted sodium bicarb will effect the experiment.

Generally, when we try to use a reaction to produce something, we try to use one reactant in excess.  That insures the reaction goes forward rapidly, and consumes all of the other one.  Since acetic acid is volatile, you could add it until the fizzing stops, then add more.  And then a little more.  If you evaporate it at this point, you can be more sure the bicarb is gone, and the acetic acid is gone too, after evaporation.  Also, you can recrystallize, but since sodium acetate is so soluble, that can be difficult.  Unless its less soluble in alcohol, or something else easily found. 

Still, a very good job on your part working out the stoichiometry.  Its sometimes very hard to get a new poster to care.  You deserve one just for that.

[yyttr3]

I don't have any proper measuring devices so i'm using cups as a unit of measure, sorry about that. I tried to measure 3 cups as precisely as possible with some error i'm sure. I tried looking up the msds for Great Value distilled white vinegar but I can't seem to find it so i'm hoping the bottle is correct in saying that it is 5% acetic acid. I'm using arm & hammer baking soda which is advertised as 100% sodium bicarb and the msds says the same. I'm sure there are some impurities but i'm hoping they're insignificant.

I don't want to use sodium bicarb in excess because i'm going to boil all of the water away and I want to have    very pure sodium acetate crystals if possible.

Forgive me for the significant figures, i've never really done anything with chemistry so I did not think of the error.

Also I have one more question, wikipedia says that sodium acetate (trihydrate) has a boiling point of 122 degrees Celsius. Correct me if i'm wrong but does that mean that solid sodium acetate dissolved in water has a much lower boiling point of 122 degrees? If I don't have a heater to accurately control the temperature when boiling off the water would it effect my results?

[Arkcon]

Right, so, like I said, use an excess of vinegar, that will be all gone after boiling and crystallizing, and there will be less chance of residual sodium bicarbonate.

The temperatures quoted in Wikipedia are for pure substances, they don't mean anything to a water solution.  I'm pretty surprised you can melt solid sodium acetate without it decomposing

[magician4]

The temperatures quoted in Wikipedia are for pure substances, they don't mean anything to a water solution.  I'm pretty surprised you can melt solid sodium acetate without it decomposing

it was the trihydrate he mentioned; and that, you can easily melt without decomposition  (57°C , if memory serves)

the "boiling point" of the trihydrate is in fact the decomposition point of this hydrate

melting point of the pure sodiumacetate is approx. 300°C , with decomposition occurring shortly thereafter (324°C, dec.)

regards

Ingo