UC

Glass stirring rods may not be anywhere near the most expensive equipment the average amateur chemist owns, but they are nonetheless ubiquitous. I also think it's nice to have at least one piece of equipment that you made in some way.

Current ebay prices put the typical 6mm wide, 12 inch (305mm) long stirring rod between $0.75 and $1 (USD) each before shipping. Instead of spending my money on a dozen or so stirring rods, I purchased a large quantity of 6mm diameter clear borosilicate rod for glass working. I won 5lbs of rods for only $5 (which comes to 51 pieces that are 24 inches (610mm) long) and picked them up since the seller was fairly local. Stirring rods that I make from this glass cost me slightly under $0.05 each. They look like this as obtained, with crude broken ends:

Normally, glass working (especially of borosilicate) is done with high-temperature torches, but for the simple application of rounding the end of the rods, a plain propane plumbing torch is sufficient. To round the ends, hold the rod vertically in the flame like this:

The glass does not soften too fast with such a cool flame, but after a minute or so, it has liquefied enough to ball up. After cooling, it looks like this:

Repeat on the other end of the rod. Of course, these rods are 24 inches long, which would make for a very inconvenient stirring rod. Half of this length is, however, a very nice size. Locate the center of the cooled rod and make a mark. You could just use brute force to break the rod, but if you have a fine metal file, you can do better. Place the rod against a hard surface and slowly "saw" with the edge of the file perpendicular to the glass rod until you make a small notch in the glass like this:

Slowly rotate the rod while continuing to "saw," using the notch as a guide until you make it all the way around. It will look like this when you finish:

I recommend wrapping a towel around the rod before you break it to prevent any glass chips from going flying. Apply a moderate amount of force to the spot where you scored the glass and it will give way with a more or less clean break. Mine, shown here, didn't want to cooperate with a perfect, flat break, but you can see that the break propagated mostly. along the scored line:

Repeat the procedure with the torch to round off the new ends. Repeat and you have yourself a lifetime supply of stirring rods for pennies.

Making a glass spatula

The typical laboratory spatula is stainless steel, which is acceptable for most commonly encountered chemicals, but fails for things like copper (II) chloride or iodine when even traces of moisture are present. Plastic spatulas are also available, but for 5 cents, It's easy enough to make one out of glass.

For this, you'll need a pair of cheap pliers without teeth, which have a reasonably large surface area to the jaws. I use a pair of 8" linesman's pliers that I took a shop grinder to. Heat up the end inch or so of glass until it starts to get soft, then insert the pliers into the flame (hence the need for cheap, as this will ruin the temper) and use them to flatten the glass, reheating between each squeeze. When working with such a low temperature flame, removing the glass item causes it to harden almost immediately. The metal on the pliers is a conductor of heat, and thus a bad choice for glass working, but still usable if you take care to reheat the piece often. Use the jaws to reshape the spatula as you please and put a bend in the end of it like this:

I wouldn't call them pretty, but they will do their job with the more aggressive solids in your chemical library. The small amount of rust on the right spatula was from the pliers and is embedded in the glass, where it can do no harm.

Theo- Heptane is a very good solvent for iodine and less nasty than benzene. You'll find that the volatility of iodine makes separation nearly impossible though. You're better off slowly adding a sulfite solution to the tincture just until it goes clear, neutralizing with NaHCO3 solution (since the reduction generates HI), boiling off the alcohol, and then oxidizing the iodide back to iodine with chlorine generated in-situ. This part is actually what ope was trying to suggest above, but I've removed ethanol from the equation. Acidify the liquid with HCl and use a dropper to add bleach with swirling until you don't get any further precipitate.

The standard lab test for phosphate is to use ammonium heptamolybdate solution. On addition to either phosphate or arsenate, an intense yellow coloration or precipitate is developed. This is due to the formation of complex phosphomolybdic acids.

Barium nitrate or chloride will give a white precipitate for both phosphate and sulfate. Once you know it must be one of the two, use the molybdate to test further.

Theo has the others correct. Bicarbonate is very hard to distinguish from carbonate, although, you should be able to add a very small amount of calcium chloride solution into the bicarbonate solution without precipitate. Upon heating or standing, you'd get familiar calcium carbonate. Calcium and magnesium bicarbonate are what make hard water "hard."

john do you do everything by the book or how about growing a brain

even if it doesn't work (idk maybe wrong hydrogen will be replaced) its logical.........

You clearly have no idea at all what you're talking about. Kindy shut it.

To the author of the thread:

The way to approach this kind of problem is to backtrack through simple reactions. The molecule you want to make has bromines on adjacent carbons. This is the result of the textbook addition reaction of bromine across a ___________.

There are quite a few ways to make a _________, but there is probably one you have learned which utilizes triphenylphosphine.

I'll let you run with the rest of the problem. Hopefully that nudge got you going in the right direction

uc's got the right idea... it won't kill you.

 

plus i'm sure chlorine is absolutely safe

i mean salt has chlorine but you eat salt, right?

No, salt has chloride anion. Chlorine gas was used to kill tons of people in WW1. At low concentrations, however, it acts merely as an irritant. Plus NaCl is pretty toxic, all things considered. It would just be hard to eat enough to kill you because of the taste.

When I was less careful with my experiments, I managed to almost gas myself once or twice. The smell clings in your nose for hours and I was coughing for quite a while.

Unless you're stupid enough to add acid to the bleach, you won't reach levels like I inhaled.

Teflon is PTFE or polytetrafluoroethylene.

Have you considered just ampouling the samples? with a little practice, this can be done somewhat crudely with a plain old borosilicate test tube and a propane plumbing torch. The halogens won't permeate solid glass. To do this with iodine, just keep the flame away from the iodine. For bromine, you should submerge the bottom of the tube in ice-salt brine to minimize volatility.

You can buy teflon sheeting and cut out a cap liner to stick in a tube with a screwcap. Tightening the cap converts this into an inert compression fitting. It will very slowly leak over time, but is readily openable if necessary.

If you don't need a shelf stable sample, seal the iodine in a HDPE jar with a tight fitting lid, place it inside a second jar containing sodium sulfite and sodium bicarbonate to absorb escaping iodine, and place the whole shebang in the freezer to keep volatility to an absolute minimum. Bromine is also best stored in the freezer inside a teflon-capped tube with the sulfite/bicarbonate secondary container if not ampouled.

Alternatively, you can buy some FEP or PFA plastic bottles, which are also fluoropolymers and entirely resilient to attack by the halogens.

http://macro.lsu.edu/HowTo/solvents/toluene.htm

For general chemistry, truly dry toluene is not necessary, but specialized reactions require truly anhydrous conditions. I believe refluxing over sodium with benzophenone as an "indicator" is the standard drying method.

It's a lovely mix of chlorine and chlorine oxides. Nothing you really want to be breathing on purpose, but a little won't kill you.

Please haelp me to make my projects on the topic to test the presence of Ni2+ in Chocolate please help me..........

What kind of levels are you expecting? I suspect that you will need a very sensitive test for this project.

Samples should probably be digested in Kjeldahl flasks with nitric and if available, perchloric acids under extended heating, before quantitative transfer to a volumetric flask and analysis.

I do not know much chemistry, or even if its some reaction happening... But the very coolest thing that I've done with thermite only uses about 100 grams of Fe3O4 + Al mix, poured into a brick hole, topped with pennies throughout the mixture. The pennies cause explosions like eruptions from a volcano! It's pretty damn cool to do and I highly recommend it (I'm not responsible for burned property or personal damages or loss of anything caused by you trying this) lol...

 

I really need a picture to show you, but it's really amazing and I'd love to know what the hell causes this violent eruption!

Modern pennies are zinc, which has a low boiling point. You can probably figure out the rest. Use some pre-1982 pennies. I bet it doesn't work (they're solid copper).

This is pretty ridiculously dangerous, what with flying thermite and zinc and zinc oxide fumes, which can cause metal fume fever.

lol. this post is priceless. can I frame it and put it on the wall so I can laugh every day?

Hello:

 

How does an atom determine whether it will form a covalent bond or an ionic bond? Does electronegativity have anything to do with it?

 

Thanks for your help in advance!

Electronegativity has everything to do with it. As you may have guessed, covalent and ionic are not absolute terms. It is a range. The bigger the separation in electronegativities, the more polar they become, until the bonds become so polar that they can disassociate.

The general rule is metal + nonmetal = ionic, nonmetal + nonmetal = covalent, but this isn't always the case. If you look, I think you'll find that a Pb-C bond is less polar than the C-H bonds in hydrocarbons, which we often think of as the prime example of being nonpolar.

If you take metals to mean the alkali and alkaline earth metals, than yes, the rule applies.

Given the summer time I've always wondered if there were any substances that would help in combating heat and humidity without using airconditioners and electrical appliances.

 

I was looking at a chemical approach. Is there something I can make that is first cheap or moderately cheap that would take away (absorb) humidity? and is there any safe reaction I could do that would have a decent endothermic reaction that would effectively cool down a small room?'

 

I think if there was I it would be well know, but I figure I would give it a shot.

 

Any help would be cool. ( that pun is for phi)

An endothermic reaction? No. Electricity is really cheap. Chemicals are not.

Dehumidifying, however, is fairly common. But, if the room is open to the outside, they're useless. They sell containers of slightly hydrated CaCl2, which absorb water given some time. They are meant for closets and basements where you might get mold or mildew. However, there are also electronic dehumidifers which cost more to get initially, but last a long time and are cheap compared to chemicals in the long run.

Personally, I bake the CaCl2 chips until they let off no more water, and use them as dessicant in a dessicator. With maybe 3L of air to operate on, they do their job in a hurry.

why exactly is this in chemistry?

Also, would this not be possible?

 

CaSO4(s) + 2 H2O(l) -------> Ca(OH)2(aq) + H2SO4(aq)

or

CaSO4(s) + H2O(l) + CO2(g) ------> CaCO3(s) + H2SO4(aq)

???

This little bit leads me to believe that you copied all the rest of that text from somewhere, because this shows next to no knowledge of chemistry.

Ah yes. At least some of it is from here: http://www.cavemanchemistry.com/oldcave/projects/acid/

post reported.

Preserved for posterity:

The manufacture of sulfuric acid presents us with an interesting lesson in industrial economics. We have seen that the roasting of sulfide ores produces sulfur dioxide as a waste product. For example:

2 PbS(s) + 3 O2(g) -----> 2 PbO(s) + 2 SO2(g)

 

 

From the beginning of metal smelting to the mid 18th Century, sulfur dioxide was simply sent up a chimney into the atmosphere. Over long periods of time, the sulfur dioxide slowly reacts with oxygen and water in the atmosphere producing sulfuric acid:

2 SO2(g) + O2(g) -----> 2 SO3(g)

SO3(g) + H2O(l) -----> H2SO4(l)

 

 

This is one important source of acid rain. Consequently, a smelter was not the ideal place to build your dream home. But the world was big in those days, the wealthy simply didn't live near a smelter, and the environmental lobby was nonexistent.

 

The discovery that indigo could be used to dye wool changed the situation dramatically. Now there was a demand for sulfuric acid but no way to produce it cheaply in the quantities demanded by the textile industry. In 1746 John Roebuck developed the lead chamber process for the manufacture of sulfuric acid. Prior to this time, sulfuric acid had been produced in glass bottles several pounds at a time. But the lead chamber process could produce sulfuric acid by the ton.

 

In the lead chamber process, sulfur and potassium nitrate are ignited in a room lined with lead foil. Potassium nitrate, or saltpeter is an oxidizing agent which we have seen when we discussed explosives. The saltpeter oxidizes the sulfur to sulfur trioxide according to the reaction:

 

6 KNO3(s) + 7 S(s) -----> 3 K2S + 6 NO(g) + 4 SO3(g)

 

The floor of the room was covered with water. When the sulfur trioxide reacted with the water, sulfuric acid was produced:

 

SO3(g) + H2O(l) -----> H2SO4(aq)

 

Notice that this process depends on cheap supplies of saltpeter and produced yet another air pollutant, nitrogen monoxide. Thus we have to pay for a nitrogen source, saltpeter, and all of the nitrogen winds up going up the smokestack. Saltpeter could be replaced with its less expensive cousin, sodium nitrate ("Chile saltpeter") but nevetheless you wind up paying for nitrogen which winds up as a waste product. Saltpeter was a significant expense.

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Merged post follows:

Consecutive posts merged

Joseph Gay-Lussac invented a process for recovering the nitrogen in nitrogen monoxide and recycling it to replace the saltpeter as a source of nitrogen to generate sulfuric acid also:

 

 

4 NO(g) + O2(g) + 2 H2O(l) -----> 4 HNO2(l)

4 HNO2(l) + 2 SO2(g) -----> 2 H2SO4(aq) + 4 NO(g)

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Merged post follows:

Consecutive posts merged

Also, would this not be possible?

 

CaSO4(s) + 2 H2O(l) -------> Ca(OH)2(aq) + H2SO4(aq)

 

 

or

 

CaSO4(s) + H2O(l) + CO2(g) ------> CaCO3(s) + H2SO4(aq)

 

 

???

 

 

 

 

This may also be of interest:

 

You can replicate the most primitive production of sulfuric acid from sulfur and saltpeter. The reaction will take place in a test tube with very little air, so we must count on the saltpeter to supply all of the oxygen. Begin by balancing the reaction:

KNO3(s) + S(s) -----> K2S + N2(g) + SO3(g)

or,

H2O + KNO3(s) + S(s) -----> K2S + N2(g) + H2SO4(l)

 

Then use stoichiometry to calculate the number of grams of saltpeter needed to react with 1.0 g of sulfur. You will bring these calculations with you to the lab when you are ready to make sulfuric acid. You will also need a 2 L soda bottle to take the place of the original lead chamber.

 

Begin by weighing 1.0 g of sulfur and your calculated weight of saltpeter and place this mixture into a clean, dry, test tube. Close the test tube with a rubber stopper fitted with a piece of glass tubing. Rinse your 2 L bottle with water and drain it, leaving the walls of the bottle wet. Put the glass tubing into the bottle and hold the test tube over a Bunsen burner. The sulfur and saltpeter will react producing a mixture of nitrogen monoxide (brown gas) and sulfur dioxide (clear gas). These react further in the presence of water to produce nitrogen and sulfur trioxide, which dissolves in the water to produce sulfuric acid.

 

What you will see initially is a violent reaction as the sulfur and saltpeter are heated. Billowing clouds of gas will be produced and the material may even catch fire. This is nothing to worry about. The gas that is produced will be brown in color and will go out through the glass tubing and fill the 2 L bottle. The bottle will seem to be filled with brown, hazy mist: a mixture of nitrogen oxides and sulfur dioxide. The gas will turn clear and trasparent over the course of several minutes as the nitrogen oxides oxidize the sulfur dioxide to sulfur trioxide. Sulfur trioxide dissoves in water to produce sulfuric acid.

Borazine is called inorganic benzene this is due to isostructural with benzene.

Inorganic benzene is more reactive than benzene. It reacts with hydrogen chloride in an addition reaction. If borazine were truly aromatic like benzene this reaction would not occur without a Lewis acid catalyst.

Benzene does not react with HCl even with a lewis acid catalyst.

Borazine is at best weakly "aromatic" due to lone pair delocalization. Even with the resonance structures, at any given time, half the ring is electron deficient molecules and the other half are electron rich, contributing to it's high reactivity.

i think sodium fluorescein is typically used for applications like this.

I joined this forum just to post my theorem (or is it theorum)

Neither. At best, it's a hypothesis. The common definition of theory would apply to this, but as this is in a scientific context, theory or theorem means that you have a large body of supporting evidence and it is a generally accepted principle.

I ended up just buying a premixed thermite, because it was actually cheaper than the separate components.

 

I prefer magnesium ribbon because it avoids any confusion to the students. They might imagine that the potassium permanganate and glycerin are somehow involved in the reaction. Plus of course the ribbon adds to the drama, makes it look a lot like a fuse on a bomb.

KMnO4/glycerin has a predictable time delay and is more consistent though. I used sparklers as my fuses and they worked quite nicely.

<Capn_Refsmmat> What's next? Tactical speedos?

<UnintentionalChaos> aka: man-thong

<ydoaPs> UnintentionalChaos, that's called a banana hammock

<mooZzZzzz> ydoaPs: only if it's big enough.

<mooZzZzzz> aka: you wish.

<mooZzZzzz> night!!!!! <runs off>

Many reactions have an "activation energy" which is supplied by heat. This allows unstable, high-energy intermediates to form, which then collapse into a lower energy final product.

The colder a reaction mixture gets, the less the molecules move around and the less likely they are to run into the other reactant(s). This is why reactions are usually carried out in the liquid or gas phase. High mobility and easy mixing of the reactants. Solids would require providing intimate contact on a microscopic level, which can be realized for some reactions by grinding the reactants together.

either that or Frank.

You sir, have my undying support in this matter.

i think my irony meter just exploded

Mine definetly did. *gets a mop*

Good thing they're not mercury-filled anymore.

new researches have shown that the temp causes increase in CO2, not CO2 causes temp rise!!!

and another fact is -rise in temp is followed by rise in CO2 by 100 years!!

Peer reviewed paper or it didn't happen.

http://www.chemguide.co.uk/organicprops/amides/hydrolysis.html

Urea is just carbonic acid diamide.