Theophrastus

totally unrelated, but I just realised the answer to my previous problem in refining the magnesium oxide, to pure magnesium. If you are to heat the magnesium oxide with charcoal, the following reaction would ensue: C + 2MgO > 2Mg + CO2 I really can't believe I didn't think of this earlier, having recently read some of isaac azimov's nonfiction books, regarding the history of chemistry. This was the primary method used by alchemists and metallurgists, to extract pure metals, from their respective oxide ores. It mentions it, in the context of copper, iron and tin, however, I can't see why it couldn't work for this. Best of luck in your electrolysis though. Tell us whether it was pure magnesium or magnesium hydroxide that formed. If the first, than your goal has been accomplished, if not, you can simply dehydrate the hydroxide to oxide, and then, do the above. Best wishes, ,Theophrastus

Sorry, simply felt a little under the weather, and decided to look through any threads that were either unanswered, or seemed interesting. Yes, I know, I must find something to do... wow, I feel like a six year old! (No reasoning behind it, I simply do) (the following has absolutely no relevance to the thread in question whatsoever, does it?) Ah well...

Of the first, yes, it is true that vinegar and acetic acid react, to produce hydrogen, and magnesium acetate, however, the time I placed it into the acid, was relatively short, and as such, the contamination would only be minor. And secondly, if what you say is true, there is a serious flaw to my theory, as magnesium, being less reactive, will not displace NaOH. Even though, this reaction, theoretically occurs, due to the magnesium ion's "excited" state, Na, being far more reactive, will following the creation of Magnesium Hydroxide, displace it. Also, I doubt a large amount of NaOH formed, as, I checked the PH of the solution itself, following my experiment, to find its PH to be between 7, and 8. Just in case, following the filtration process, I'll boil down the solution, and see, whether a change occurs.

I've conducted my PH test, this evening, reaching the conclusion, that it was in fact magnesium hydroxide, layered upon the magnesium electrode. Further proof of a reaction, was after the completion of the titration, I poured out my indicator solution, getting rid of any remaining salt, upon my metal, with acetic acid, I found that all across the magnesium anode, there were tiny pockmarks, where the ionized magnesium, reacted with the water, forming magnesium hydroxide. Once the salt was neutralized, the magnesium, lost in the reaction, was evident. Good idea, with the gases! (the collection of them, I mean) I could give it a go this weekend, if I have enough free time on my hands. Thanks for the corrspondence to my question. I think, that if I'm unable to figure out how to separate the two precipitates, by the week's end, I'll simply filter both out, and start some testing. Flame tests, seeing if they react with particular compounds, if I get any clues towards what they may be. If I find out, I'll post my resolve, on this forum. Ps: Sadly, I do not live in arizona, or even USA. Think further north, up in Canada. Its really quite a shame that there are so many limitations, in obtaining chemicals up here. Cheers!

Excellent idea, and you're right that it does seem strange that a white powder is laden upon the "anode." My only explanation is that either it is salt from the solution, or else magnesium hydroxide, as the excited magnesium might react with the water: Mg* + H2O > MgOH + O. I'll do a PH test tonight, if I can, then share my results tomorrow.

As everyone has mentioned, the energy required to produce the desired effect, would be huge, and this is the ultimate obstacle. You see, while it is quite easy to think of this from an outside perspective, in thinking of what's right, business and industry looks towards what is profitable. This is why peoples' very lives are exploited in developing countries, by companies. This is why certain types of green technology that is possible is not implemented. These concerns, to a certain degree, are understandable. However, unless there is a way that is profitable, a use for the carbon dioxide, that is then captured, I seriously doubt any large scale developments in the future. I've posted a link below, showing an interesting method of transmitting CO2, a by- product of fossil fuel and coal burning, to other industrial uses. This is why the recycling of metals has become commonplace so quickly, because, the process of recycling these metals, is far more energy efficient, that simply extracting them from their ores. The point I stress is that business, generally speaking, will lean towards what is profitable, and what is economic. And in respite of that rather melancholly note, I have a lovely quote for martin luther king: "Cowardice asks the question: Is it safe? Expediency asks the question: Is it politic? Vanity asks the question: Is it popular? But conscience asks the question, is it right? And there comes a time when one must take a position that is neither safe, nor politic, nor popular, but because conscience tells us it is right."

In regards to what has been said, firstly, one has to become careful of the electrodes you use, is this is how you're planning to attain magnesium, as while magnesium will precipitate into the solution, most metal anodes, will then react with the sulfate, for example, if you are using copper electrodes, the reaction would form magnesium and copper sulfate. The magnesium, now magnesium hydroxide, having been dissolved in water, would then displace the copper, and you would be left with copper hydroxide (which actually has quite a number of useful applications!) and magnesium sulfate. To summarise, using any metal, that is less reactive than magnesium, would simply lead to the reverse of the initial reaction, going nowhere. Due to this, I recommend you use carbon rod electrodes. For this, graphite will do. However, I have a feeling that the magnesium, would then become magnesium hydroxide, reacting with the water. But hey, you can't knock it, 'til you've tried it, right? Best of luck! ,theophrastus ps: let us know the outcome. It may be quite interesting! (magnesium powder, could be quite useful)

I think I have realised a minor flaw in your process. It certainly seems rational that if alkali metals react with water, to isolate them, would require the use of a nonreactive, nonpolar solution, however, in the case of solutes and solvents, like dissolves like, and as such, the ionic solute, which would maximize the electrolytic properties of a solution, would be unable to dissolve,in the solution, and thus, the electrolysis would be deemed ineffective. Sorry to spoil the fun. (though let me know if I'm wrong; alkali metals, are rather fun!) Best wishes!

I'm unsure about your solar cell, but I think I can hazard a guess at what occured. In the electrolysis of water (Or any substance, for that matter), there is no load, on the electrical circuit, and as such, what can occur over time, is the "burning out" of your battery, due to overheating. It's similar to what occurs, when you take a battery, and connect a wire, attaching the battery's negative and positive end (do not actually try this, (for safety reasons), simply take my word for it). With no load in between, all the electricity does, is heats the battery and the wire. Eventually, this overheating will lead the battery to burn out. Certainly this is an extreme example, as the electrolyte, not bearing the same aptitude of conductive properties, as a wire, offers some manner of load, (well, not a load, rather an obstacle, if you will, for the electricity). However, the purpose of electrolysis, is rather for this current of electricity, flowing through the electrolyte, to produce a negative charge on one electrode, and a positive charge on the opposite. This charge, will then induce the dissassociation of ionic, and polar compounds, such as (theoretically, not practically speaking) water, salt, etc. By "breaking apart" these compounds, you can obtain, their components. (ie NaCl > Na + Cl, 2H20 > 2H2 + O2, etc.) As for whether the process will damage your solar cell, that I am unsure of, as I'm not very well familiar with the basis of the photoelectric effect. Hope this helps! ,Theophrastus

My apologies, for ommitting the fact, but I am confident that I used a magnesium anode, and a copper cathode for the experiment. Any thoughts? Thanks, for the input, though! Quite interesting!

To Salter: I doubt that the black chemical is magnesium oxide, due to the fact that magnesium oxide is white, however, the idea you gave me about the copper chloride, was certainly interesting, as I think that you were right in saying that the yellow colour of the solution was due to CuCl2, and I think the green colour, I noticed, earlier on, was CuCl. However, I doubt that my precipitate was CuCl or CuCl2, as most chlorides, (with the rare exceptions of Pb 2+, Ag+, Hg2 2+, etc.) are in fact soluble. Thanks for the earlier not though; copper(II) chloride has its uses, and I suppose I could extract it, were I to perform a similar experiment again, however, it seems I've sadly returned to the start. (Insert frowning face) I still believe that the black is copper(II) oxide. (correct me if I'm wrong.) (Sudden thought!) Hold on... couldn't one of the precipitates be a chlorine- oxygen compound, like a chlorate, hypochlorite, etc. No, but since I was using copper in its usual valency (II), wouldn't it be dichlorate, dichlorite. Personally, I know little about these compounds. Also, none of the precipitates can be magnesium, as aren't most magnesium compounds white? I also examined the magnesium anode, more carefully, and decided to finally note a thin layer of crusty, white powder upon it. Examining it thoroughly with a magnifying glass, I also found miniscule traces of a brownish gold residue. Could this perhaps be traces of my upper precipitate, that was left inscribed upon the anode? Any thoughts?!?

Oh my god! Pyrophoric? At first, at your comment, I admit I was greatly dismayed, as looking it up, I did find, to my surprise that it was pyrophoric, however, my specimen, does not spontaneously burst into flame, which is a good thing, probably a testament to its relative purity, perhaps, however I seem to doubt this fact. The process described on a site, of the compound's burning to produce either sulpher dioxide, or, if no oxygen is available, its decomposition to iron, and sulphur, I found to be cohesive with the process I envisioned, in regards to the reaction itself, but far more dangerous, than I thought, as it seems that while the reaction would occur, it would seem to accelerate violently. Certainly not something for the faint of heart, and those who are not yet experienced enough, to properly contain such a reaction. Truthfully, I fall into the second category. That taken into account, I've disbanded my original idea, and instead, I suppose I can use UC's idea, of heating metabisulfites to produce a metal oxide, and sulphur dioxide, creating a simple ester, within the apparatus to check for any leaks, prior to performing the actual experiment.

I would have suggested a single displacement reaction, however, magnesium is quite reactive, and the only more reactive metals, I can think of, are the alkali metals, as well as some of the alkali earth metals. an easy way to obtain magnesium, is through the body of pencil sharpeners, as those are either made from steel or magnesium, and from that point on, you can simply judge it by the weight. The blade, however, is always made of steel, so its best you remove it with a screwdriver. In regards to chemically producing it, with epsom salt, you can react the epsom salt, with ammonium hydroxide (ammonia + water solution), producing ammonium sulfate and magnesium hydroxide. (Yes, I know, same subtance as would be procured in your electrolysis, but you would obtain a better yield), which when heated, produces magnesium oxide... and then.... I have absolutely no idea, what to do. Best of luck!

Such as? (He says with a hopeful tone)

To Rob: Ah, now I see. Theoretically speaking, it should work, however, I've tried the experiment myself, and found the yield, to be most unsatisfying. However, I did perform the reaction over a span of two days, and i suppose, were one to let it ensue longer, the numbers may add up. I was thinking that one could theoretically increase the yield, by using a catalyst, adding heat, (something along those lines) etc. However, perhaps, you may outdue me, as I do not pretend to be an expert. And I agree with UC in looking up the electrodes based on your needs and setup, on google. (Ah google, the ultimatum of search engines) Cheers! Best of luck in your experimentation! To Hermanntrude: I sense I did something wrong, but what? If it was the degrees Kelvin, is there perhaps a problem with using that system of measurement, or is it that the number I gave, was off by one. Looking back, I believe it is so... or... what?

Thinking, it would better my opportunity for an answer, I decided to continue, one of my questions, here, as it seems vaguely relevant to this post, which while old, at a time was discussing the electrolysis of copper, and various possible products and bi- products that could be made as a result. 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 black- brown, while the upper one a dull, pale orange. What could have gone wrong? What are these precipitates? Can they be of any use? Help with these questions would be most appreciated! I would have dried the precipitates, and made a photograph, to show, however, having two distinctly different precipitates, seems to complicate matters. How can I remove these two precipitates from the solution separately, withy no knowledge of their chemical formula? ,Theophrastus ps: In retrospect, I believe that the black substance, may be copper (II) oxide, from the anode (2Cu(II) + O2 > 2Cu(II) O), however I am unsure, and uncertain. Any ideas???

Recently, I've been thinking a lot about preparation, and methods of synthesis for various compounds, and amongst these, I've thought of producing sulfur dioxide, in some way, and then bubbling it, through a dilute hydrogen peroxide solution. As a result, this ought to yield a dilute solution of sulfuric acid, (about 5%, based upon the type of hydrogen peroxide I have) which I can then boil down, to a more feasible concentration of approximately 15%. (roughly) My only question, is in the generation of the sulfur dioxide. The obvious way, would be simply to burn sulfur crystals, however, I am unsure where these can be obtained. I do have various sulfides, (FeS2), which I suppose, upon heating, will oxidise to SO2 and Iron Oxide, but what are the temperatures required, for a such a reaction, to take place? If the temperatures are too great, are there any alternative ways to produce the SO2, necessary for the reaction to occur? Ps: To prevent any furor, or unnecessary worry, I am well aware of the toxicity of this compound, and shall plan accordingly.

Thank you, I was simply making sure, however, I was thinking, what is the precise decomposition temperature for NH4Cl, because I was thinking, that if there is enough of a range between the boiling points of its two components, ammonia, and HCl, I suppose I could funnel the ammonia, into water, where it would dissolve, forming an ammonia solution, (hopefully) leaving the HCl, in the previous flask. Something along the lines of a basic distillation. Any thoughts?

In regards to experimental uncertainty, I have always thought of it in the context of measurement, as we are all human, it is only human nature for our measurements to be correct, to a relative degree. Scientists often like to take note of the theoretical uncertainty of a certain measurement. For example, (doing so with a completely open mind) use a 30cm ruler, to measure the circumference of a door, to the nearest mm. Repeating this experiment several times, one will find that the various measurements shall not be absolutely coherent with one another, but shall be, perhaps, between one and two cm's long or short of one another. Thus experimental uncertainty. In regards to the constant mention of the work of Heisenberg, I disagree in throwing him into this arguement, as he proved that at a quantum mechanical level, it is impossible to make accurate measurements of speed, length, distance, etc., he was not refering to the macroscopic world, however, I admit that they may bear implications regarding measurement, that transcend his actual context.

I truly doubt that there is any way to get acetic acid, a carboxylic acid, to form an alkyne, as generally, in organic chemistry, there are patterns to these sorts of syntheses, such as oxidizing alcohols to form carboxylic acids, or esters by reacting a carboxylic acid with an alcohol. However, if your question was more general, in simply creating acetylene, I must forewarn you that acetylene is incredibly volatile, and combusts quickly and violently. I've posted a link below, a testament to the danger this compound poses. A question one must ask is, why do you want to produce it anyway? (Meaning no offense, but... safety first.)

Thank you for the most needed information, however, upon decomposition, would not the products, simply react with one another to form ammonium chloride again? Is there perhaps any way of isolating the individual chemicals?

I certainly understand that einstein, being human, made mistakes, however, I see little point in reading a book all about his various mistakes, embellishing scientific failiure. Personally, I find such a thing, very degrading. One interesting thing, is that his "greatest mistake" (He even acknowledged this himself) , the cosmological constant, which he added to his relativity equations, in an attempt to retain the uniformity of Newton's static universe, was not in fact so great a mistake. It was certainly foolish, in that, had it simply been left out, he might have been able to predict that the universe is either expanding or contracting, however, science has recently found that there may be a smaller (certainly not allowing for a static universe, in its weak influence) cosmological constant, built in to the fabric of spacetime. It's thoughts like these that simply make me feel happy inside. I think the word we're looking for is serendipity.

With a little thought, I made a sudden epiphany recently, in noticing that ammonium chloride (NH4Cl), can actually be split into two separate chemicals, NH3, and HCl. I thought of various methods, how this can be done, however, in ionising, the salt does not form the compounds above, but rather ammonium ion (NH4+) and chlorine ion (Cl-). Is there any way perhaps, to coerce this salt to decompose in the compounds described above, and how can this be done, without having the two products react, to form ammonia salt again. Any ideas?

An interesting thought, has recently occured to me, an idea of how to synthesize ammonia, by means of two salts I had in my possession, Ammonium Chloride, and Calcium Hydroxide, my proposition, is to create ammonia, by heating the two salts to the necessary degree, and thus, a reaction like the one below would ensue: 2NH4Cl + Ca(OH)2 > 2NH3 + CaCl2 + H2O I know that ammonia has a high solubility in water, and thus, I could perform the reaction in an erlenmeyer flask, with a holey stopper, where a glass tube would be inserted, and the ammonia, thus bubbled into water, forming NH4OH (Ammonia dissolved in water) However I am confused. Is a reaction like this really ethical? And if so, what degree of heating (activation energy) is required? Should the reaction be emersed in an aqueous solution, or can it be performed without this? And finally, if the setup above is proved impractical, is there any other way to safely generate ammonia, using basic chemicals?

Expanding upon what gonelli said, if you are doing an electrolysis of NaCl, you are to be forewarned that the melting point of NaCl is 1074 degrees Kelvin (800 degrees celsius/centigrate). The heat required, is what makes this process difficult to enact. If you were refering to the electrolysis of water, adding salt, to increase its electrical conductivity, certainly, various metals, could be used. Personally, I like to use copper and magnesium, however, I know that many like to use graphite. Some also use nickel or platinum, due to their catalytic properties, accelerating the reaction. To properly answer this question, I must ask, are you doing an electrolysis of NaCl, or H2O, and what do you desire to achieve or produce, from this reaction?