Theophrastus

While it certainly isn't a lab grade quality, eggshells are approximately 95% calcium carbonate, so if one were to powder the shells, by means of a mortar and pestle, being careful to peel off the inner membrane first, one should come up with calcium carbonate powder, of a relative purity. I've tested this, by reacting the remaining solution with acetic acid, and it works quite well, however there does seem to be some organic residue remaining, a testament to its relative purity. You can also buy it as an antacid in most drugstores. While I am unsure in regards to how one could extract it from limestone, you can test for calcium carbonate by reacting your mineral extract, with a strong acid, however, while reacting less vigorously, weaker acids, such as acetic acid will suffice. If the mineral seems to bubble, then you know that there is calcium carbonate present. (example of the emergence of CO2, following the addition of a strong acid, such as HCL) CaCO3 + 2HCl > CaCl2 + CO2 + H2O

Ethanoic acid, more commonly known as acetic acid, if performed electrolysis upon, would probably break apart to form hydrogen, and acetate ions, as shown below: CH3COOH (insert reaction arrow here) H+ + CH3COO- Hypothetically speaking, the hydrogen would bubble up, from the anode, while acetate ions, would be despelled from the cathode. Then, depending on the metal in use, (generally speaking) the acetate ions would react with the cathode metal, to form that metal's acetate. For example, were one to use use copper, as a cathode, the resulting compound would be CuOOCCH3, copper acetate. In regards to your second question, the sodium hydroxide would (based upon the concentrations) would either neutralize, or be neutralized by the ethanoic acid, as shown in the reaction below. CH3COOH + NaOH (insert reaction arrow here) NaCH3COO + H2O As you can see, a neutralization reaction occured, forming sodium acetate and water. In regards to your last question, this I am unsure of, as magnesium, generally forms an outer layer of oxidation (rust) to protect the rest of the compound. (MgO) However, upon contact with moisture, this forms Mg(OH)2, as in the figure shown below: MgO + H2O > Mg(OH)2 As such, were one to use this practical application, the magnesium hydroxide, would react, as shown below: Mg(OH)2 + 2CH3COOH > Mg(CH3COO)2 + 2H2O Having destroyed the outer oxide layer, the ethanoic acid would then react with the pure magnesium, breakinjg it down, as shown below; Mg + 2CH3COOH > Mg(CH3COO)2 + H2 This process would then form magnesium acetate, and hydrogen gas. Hope this helps! ,Theophrastus

Greetings. Not long ago, I was conducting a simple electrolysis experiment, performing electrolysis of water, using a voltage of approximately 13- 15 volts, in a saturated aqueous salt solution. I chose to use a magnesium anode, and a copper cathode, to my surprise, a substance, faintly green in colouration, began to rise from one of the electrodes. Fearing this was chlorine, I quickly disbanded this setup, building a different one, in which the reaction would take place in an erlenmeyer flask, with a holed stopper. Through the hole in the stopper, I placed a curved glass tube, which would release any gases made in the reaction, into a separate vessel. The wires connecting to the cathode and anode, were also inserted into the hole. Soon, the reaction ensued, as a stream of bubbles vigorously rose from the anode. As time passed by, I quickly watched the solution change colour from a pale green- yellow, to a soft golden yellow. However, upon my return to the room, the next time, the solution had become discolored and opaque, due to the presence of a dark brown precipitate. I waited a while for the solution to settle, to find two precipitates of varying density, layered upon each other, at the bottom of the flask. The lower one was a dark brown, while the upper one a dull orange. What could have gone wrong? What are these precipitates? Can they be of any use? Help with these questions would be most appreciated! ,Theophrastus Merged post follows: Consecutive posts mergedIn retrospect, I believe that the black substance, may be copper (II) oxide, from the anode, however I am unsure, and uncertain. Any ideas??

Agreed, particularly in chemistry, safety is a concern, however, I must admit I've detected a rather accusatory tone to your comment, and as such, can say that I am not interested in "bangs," to try and avoid such a sensitive subject. In regards to common names, I have already found MgSO4 (epsom salt), and Na2Co3 (Washing Soda), and the like. My need for a way to synthesize acid arises from the fact that despite a fair amount of searching, I was unable to find muriatic acid. In regards to the rules and code of conduct, I'm grateful that you pointed this out, as it is useful to know the rules by which one is to play. Cheers!

Thank you Hermanntrude, your advice was most appreciated. I myself, have a rather dirty wine vinegar in my possession, however, I'm sure this can easily be ammended. Cheers.

In regards to the hydroxide ions, I would imagine that one could safely dispose of them by then transfering the vapours into an appropriate solution, in which they might then react to form stable, insoluble hydroxides. As for your comment regarding the heat of the compound, you are right. I was thinking, in more of a hypothetical context. Hydroxide, is an ionic compound, and as such, would be expected to have a rather high boiling point. In regards to a more practical approach, I am at a bit of a loss. Could not one intermix another compound, perhaps some other alkali or alkali earth metal salt, to decrease the boiling point. (Granted that the question is, to what degree can this be done...) However this would leave you with an alloy of the two metals. Alkali metal alloys in particular, are infamous, due to their instability. Certainly a process that is neither easy, nor safe.

I've currently been thinking of performing various reactions, using acetic acid to precipate several of the reaction products turning them to acetates. As per, doing so, requires one to actually have acetic acid on hand, so my question is, how can one produce acetic acid? I understand that vinegar is about 10% acetic acid (well, the concentration may vary...), so might there be any way to distill this from vinegar, based upon its othe components (ie water, spices, etc.)

Hello, I have currently, after doing a variety of theoretical work have begun to venture into the world of the chemical, only to find my first impediment. Obtaining chemicals. To my dissappointment, canada, seems not to be a principal nation in the amateur chemistry department, and as such am at a loss to obtain various key chemicals. My primary concern, as of now, is the obtaining of strong acid or base, such as hydrochloric acid, sulfuric acid, sodium hydroxide, etc. I have currently been able to attain magnesium, copper, iron, iron (III) oxide, Copper sulfate, Calcium Hydroxide, ammonium chloride, potassium hexacyanoferrate, sodium hydrogen sulfate, ammonium iron (III) sulfate. Thoughts and help on any sorts of chemical syntheses I might be able to accomplish with these chemicals is most appreciated. ps:(just please don't mention the obvious displacement reactions I could perform with CuSO4) pps: While help in chemistry is appreciated, if you are to find a Canadian chemical supplier, please let me know! ppps: One idea I had was the obtaining of ammonia, by reacting NH4Cl and Ca(OH)2 as such: -2NH4Cl + Ca(OH)2 - 2NH3 + 2(H2O) + CaCl2 Then reacting it with copper sulfate to form a coordination complex, after extracting it from the solution.

In regards to this predicament, I'm sure you can obtain sodium, by means of an electrolysis of sodium hydroxide, obtaining sodium metal, at the negative electrode, and hydroxide anions on the positive electrode. Hope this helps, ,Theophrastus