Man has prepared dilute acetic acid since the Agricultural Revolution by simply letting mead sour in air. Hence it is also likely that sodium acetate has accidently been produced since then in an impure form. Today, sodium acetate is used as a food preservative and in "instant" heating pads. It is easily prepared by reacting baking soda and vinegar.
NaHCO3 + CH3COOH --> CH3COONa + CO2 + H2O
Sodium acetate is actually a suprisingly useful chemical in the amateur labortory, and this usefulness is magnified by its ease in preparation. Distilling sodium acetate with sulfuric acid gives concentrated acetic acid. It nearly gives glacial acetic acid, but since some water always comes over it is necessary to suck this up with some drying agent to convert it to the glacial form.
Salts of carboxylic acids are handy for producing small amounts of hydrocarbons. Upon heating with lime or lye, sodium acetate yields methane while electrolysis of a concentrated solution gives ethane.
Another reason I want to synthesize sodium acetate is because you can perform neat supersaturation experiments with it! Supersaturated solutions of it are relatively stable and with a seed crystal or the touch of the finger, you can immediately grow crystals of sodium acetate trihydrate right before your eyes. This process is highly exothermic and hence its application in heat packs.
40.0g baking soda was slowly added to 575mL of vinegar, or 5% v/v acetic acid. I added excess acetic acid so that it simply would evaporate upon boiling. The solution was stirred for about 10 minutes before boiling commenced. When a 200mL solution was obtained, the mixture obtained a definite golden-yellow color. I ceased heating when a bubbly tan precipitate began to form. This precipitate was dryed in a toasted over at 300F for 2 hours. Anhydrous sodium acetate was thus produced and this appeared as a fine white powder. However, when a tiny bit of water is added to some of the product, it turns back to the tan trihydrate form. The discoloration is probably due to organic impurities in the vinegar because although the vinegar is distilled, it is produced via fermentation. The final dry yield was 34.4g or 88% of the theoretical.
I was able to perform this experiment in college since it requires almost no equipment. Indeed, my friend and I used the bottoms of aluminum cans as our "beakers". 600mL of vinegar was added to about 37g baking soda. Since we did not have a scale, we estimated using baking soda's density of 2.159g/cm3 that this amounted to 3 1/2 teaspoons of sodium bicarbonate. The solution was boiled down to dryness, giving anhydrous sodium acetate as a white powder. After several failed trials, we discovered a procedure that consistently resulted in a stable supersaturated solution. Enough boiling water was added to the white powder to dissolve the product. Upon cooling the brown solution to room temperature, large clear needle-like crystals of sodium acetate trihydrate formed amongst brown goo. These crystals were isolated and heated until they melted in their own water of crystallization, giving, again, a dark brown solution. This solution, however, was stable at room temperature with respect to crystallization. A seed crystal was then dropped into the solution resulting in instant crystallization, radiating rapidly outward from the point of contact. The can heated up considerably as a result. Pretty nifty!