Sciencemadness Discussion Board - AgNO3......?

AgNO3......?

mert - 22-8-2006 at 11:24

how am i a AgNO3 solution(high pure and low pure) produce with cheap via in a laboratory....?

not_important - 22-8-2006 at 11:56

First find some cheap silver. Then tell me how I can get some.

The old standby is to dissolve silver coins in nitric acid, evaporate to dryness, then fuse the silver nitrate. Copper nitrate, and the nitrates of several other contaminating metals, is less stable to heat than silver nitrate and decomposes to the oxide. After cooling the silver nitrate is dissolved in distilled water and filtered off from the copper oxide. Repeat the operation once or twice, then crystalise the silver nitrate.

Electrolytic refining of the silver can be done as well.

You'll need pure nitric acid, check the treads on making HNO3. And go the the 'library' section, download some of the old general or inorganic chemistry books; they often tell how to make reagents.

Cesium Fluoride - 22-8-2006 at 12:09

Not sure where you at, but in the United States...Dimes made before 1965 contain 89.24% silver and 10.76% copper. Quarters made before 1965 are 90% silver and 10% copper. Modern day dimes and quarters are 91.67% copper and 8.33% nickel. Silver solder is another source.

The_Davster - 22-8-2006 at 12:12

Go down to your local coin collector store, they usually sell silver by the ounce. Cost me around 10$ CDN an ounce.

S.C. Wack - 22-8-2006 at 13:00

[rant]The price of silver has doubled over the past 2 years and may soon triple over the past 3 years...Don't buy high they say, but high might be later rather than sooner. They said gold prices wouldn't soon go back up with the Swiss dumping and all...the price has nearly tripled since then. As long as industrialization + population growth continues, demand will reduce supply of all but the most common metals like zinc(the price of which has more than doubled in less than a year)...copper has quadrupled over 4 years, new mines coming or not. You might want to buy now, even though you're nowhere near the ground floor; which might not be coming back anytime soon.[/rant]

mert - 22-8-2006 at 13:12

everybody alot thanks. i have the chemicals. so i can just do it.i am working a lab

. i will try doing it tomorrow

neutrino - 22-8-2006 at 15:35

Fusion leaves the silver salts and destroys the others? Given the instability of Ag+ ions, I would think it would be the other way around.

I made my silver nitrate by a very simple method. First dissolve the coin in nitric acid (any concentration is fine if you have heat and time). Evaporate to get crystals and wash them with ethanol until the blue color disappears. Copper nitrate is ~10x more soluble than silver nitrate in ethanol, so you're eventually left with almost pure silver nitrate crystals.

If you want to recover the silver you washed away, just throw in some copper wire or vitamin C.

garage chemist - 22-8-2006 at 16:01

Silver is commonly purified by dissolving the alloy in HNO3 and adding HCl. Copper stays in solution, and silver precipitates as AgCl.

AgCl can be purifed by dissolving in conc. NH3 solution (forms diamminesilver complex, AgCl is reprecipitated upon acidification) and filtering. This is only necessary when mercury is present (Hg(I) forms an insoluble chloride too, but this doesn't dissolve in NH3).

AgCl can be reduced to the metal sponge by hot NaOH solution + sugar (I tried it, it works very good- suspend AgCl in a bit more than equimolar amount of NaOH solution, heat to about 90°C and slowly add sugar solution while stirring vigorously).
Then redissolve in HNO3, evaporate and you have your pure AgNO3.

skippy - 22-8-2006 at 16:32

Isn't the silver ammonium complex dangerously explosive? Felt it was important to warn people!

Silver chloride will disolve in DMSO and can be precipitated by dilluting the DMSO with water. I read an abstract on it and I believe they added excess chloride ions to the DMSO via calcium chloride and that increases the solubility of the silver chloride.

not_important - 22-8-2006 at 17:54

Quote:

Originally posted by neutrino
Fusion leaves the silver salts and destroys the others? Given the instability of Ag+ ions, I would think it would be the other way around...

Empirical data trumps intuition:
Cu(NO3)2 , 3H2O MP 114.5 C, -HNO3 @ 170 C

AgNO3 MP 212 C, decomp. 444 C

Silver nitrate was melted and cast into sticks for use in cauterization, in that form it was know as lunar caustic.

If you ever owned "The Boys Big Book of Exceedingly Dangerous Science Things" or similar, you would be familiar with this method. It was common in amateur science boooks up to WW-II.

This works well for technical grade silver nitrate, especially if you evaporate the AgNO3 filtrate and do the fusion again. Use a fine filter, coffee filters are not going to cut it and likely will reduce some of the silver nitrate. Use good grade filter paper, or better fine fritted glass.

For high purity, dissolve the tech nitrate in distilled water, then precipate AgCl with pure HCl (make your own constant boil ing grade via distillation, leaves behind the iron in tech grade). Don't add too much HCl, best to calculate about how much is needed, add a bit less than that, then add more slowly until no more AgCl is formed.

After that follow garage chemist's NaOH and sugar to get pure silver or silver nitrate. This leaves behind less noble metals, some of which will survive the fusion. Precipitating the chloride from the original alloy solution tends to occlude some copper in the silver chloride.

The fusion method aworks well with several inorganic compounds. Manganese sulfate is one, it melts at 700 and decomposses around 850; the iron sulfates fall apart well before then, iron being a common impurity in technical grade manganese compounds.

YT2095 - 23-8-2006 at 01:18

a slightly different method is to dissolve your coins (or whatever) in the nitric acid, then put in thick copper to displace all the silver, filter out your silver and wash well.
then add this silver to fresh nitric acid.
I store mine as a soln with a little excess silver in the bottom of a light-tight bottle.

Engager - 24-8-2006 at 01:24

Sorry for bad english

The best way to make pure AgNO3 is to use silver alloy coins as starting material. Silver coins are always made from copper-silver alloy. First of all check out the fineness of silver in coin, it's always printed on the coin some place. For example if you see fineness .700 or just 700 that means that coin contains 70% of silver and of 30% copper. Weight the coin and calculate weights of copper and silver. Now you can proceed.

1) Dissolve the coin in 20-40% HNO3, process takes some time and may be accelerated by addition of hydrogen peroxide. After that you must transfer a part of the solution to another flask,the part in % by volume is calculated by the rule: % of copper in coin alloy is multiplied by 3. For example if coin if .900 fineness that means that 10% is copper you have to transfer 30% of solution to another flask.

2) Transfered solution is threated with KOH, that results in formation of Cu(OH)2 and Ag2O mixture, KOH is added until they stop to appear (you can also calculate the mass of KOH needed if you want, using the formula [weight of Coin in g] = Wc ; [Cu amount in wt %] = Cu% ; [Ag amount in wt %] = Ag% ; [weight of KOH] = WKOH : WKOH = ((Wc*Cu%*2*56.1/63.5) + (Wc*Ag%*56.1/107.8)) * (Cu%*3) gramm). The preticipate is filtered of and washed with large amount of water, and is added to the remaining part of starting coin solution and is well mixed. That results in preticipation of all copper part from solution in the form of insoluble Cu(OH)2, leaving the pure silver nitrate solution:

Ag2O + Cu(NO3)2 + H2O => 2AgNO3 + Cu(OH)2

You will see that blue alloy solution (because of copper cations) will turn to almost colorless. Cu(OH)2 is filtered off and discarded.

3) Pure AgNO3 solution is evaporated on open gas burner to form pure solid silver nitrate. If all the mesurments during the process are taken carefully AgNO3 product will be analytical grade.

[Edited on 24-8-2006 by Engager]

[Edited on 24-8-2006 by Engager]

[Edited on 24-8-2006 by Engager]

woelen - 24-8-2006 at 01:58

Engager, what if your initial solution is still acidic. In that case not all copper will be precipitated as Cu(OH)2. When you dissolve silver/copper alloy in HNO3, then you always need excess HNO3, otherwise not all metal dissolves.

Another point is that you will never get analytical grade chemicals from a single step process. Even if the procedure works OK, you need to recrystallize the AgNO3 in order to make it really pure.

Final point: Using a pre-computed amount of KOH is very tricky. KOH is very bad with respect to composition. Commercial high quality KOH may be less than 90% pure, the rest being water. It can also be 85% pure, the rest being water, or maybe 95%. Who knows..... KOH can be very pure, except for water content and the amount of water in it is very variable. So, yoy need to add KOH, until no more precipitate is formed. That, however, is a very cumbersome thing to do. One can also add some excess KOH.

I like the idea though of Ag2O replacing copper salts by formation of Cu(OH)2. Ag2O indeed is slightly soluble, forming Ag(+) and OH(-) in solution. Apparently this is sufficient to replace copper ions by silver ions.

not_important - 24-8-2006 at 03:11

Even attempting to titrate to an endpoint of no free copper can be tricky. And as woelen said, simply taking the solution from dissolving silver alloys and precipitating out the silver with a base metal is likely to encounter difficulties from leftover acid. You can use an excess of copper (or iron or zinc or ...), but you're just dumping more cations into solution, where they may get trapped in the silver metal being built up.

The thermal decomposition, chloride precipitation, and electrorefining all separate silver from most more and less electropositive metals. The thermal route does nothing for alkali and alkaline earth metals, so if you used sodium nitrate and sulfuric acid to dissolve you coin, the thermal method is only useful for copper removal; the sodium has to be handled some other way.

Engager - 24-8-2006 at 03:32

Quote:

Originally posted by woelen
Engager, what if your initial solution is still acidic. In that case not all copper will be precipitated as Cu(OH)2. When you dissolve silver/copper alloy in HNO3, then you always need excess HNO3, otherwise not all metal dissolves.

Agree with that, but if you regulate pH of this solution with Ag2O produced on other step all will be fine - just take slightly larger part of sol to threat with KOH.

Quote:

Another point is that you will never get analytical grade chemicals from a single step process. Even if the procedure works OK, you need to recrystallize the AgNO3 in order to make it really pure.

That's right, i think it's obvious, so i just forgot to write it. Process realy works and works well, it was taken from russian preparative chemistry book by Karjakin and Angelov "Pure Chemical Reagents".

Quote:

Final point: Using a pre-computed amount of KOH is very tricky. KOH is very bad with respect to composition. Commercial high quality KOH may be less than 90% pure, the rest being water. It can also be 85% pure, the rest being water, or maybe 95%. Who knows..... KOH can be very pure, except for water content and the amount of water in it is very variable. So, yoy need to add KOH, until no more precipitate is formed. That, however, is a very cumbersome thing to do. One can also add some excess KOH.

KOH excess doesn't matter because KOH remains in solution, and you transfer the filtered,washed solid Ag2O+Cu(OH)2 to the rest of coin solution. Solid Ag2O+Cu(OH)2 washed with water will contain no KOH.

[Edited on 24-8-2006 by Engager]

mert - 24-8-2006 at 14:13

ok.i understand.but i have a problem. i dont have silver alloy coins

what can i use silver alloy coins instead of?where can i obtain silver alloy coins? example; can it be photograph film?or other things?

[Edited on 24-8-2006 by mert]

The_Davster - 24-8-2006 at 14:15

As I said above, stores that sell coins for coin collectors.

mert - 24-8-2006 at 14:26

ohhh ok..but i will produce continuous it so i dont want buy it. i must obtain cheap it. there is silver alloy in the photography films. i want to use it or other things

Fleaker - 24-8-2006 at 17:40

The cheapest way to go is to buy silver bullion (as in ingots which have no numismatic value), melt them and cast them into shot/roll them to increase surface area, and finally dissolve in 6M nitric acid. You can then evaporate the solution to dryness and heat in a porcelain evaporating dish to expel the remaining nitric acid which will decompose at those temperatures. I have heated silver nitrate before to 'clean it up' and it does not decompose until dull red heat. I find it is the best method for removing nitric acid from an AgNO3 solution.

not_important - 24-8-2006 at 18:49

Photographic film uses pure silver; but there is so little in a bit of film that large amounts must be processed. Most if not all photoprocessing facilities reclaim silver from the used processing materials.

You don't even need to cast ingot into shot, if you're not in a rush. Silver bullion will give you tech silver nitrate with no more processing than drying. You'll still need to at least recrystallise for higher purity, more work if you want extreme purity. Bullion is best if you are making large amounts, in-circulation coins for small amounts provided your country still has silver coins; sometimes you can find silver scrape from jewelers, again might be best for small amounts but it will not be cheaper that current rates for silver and will require the processing to remove copper.

It comes down to you will be very unlikely to find a source of silver for any less than current spot prices. If you need only a bit then long searches of second hand stores and the like may turn up overlooked silver odds and ends. Beyond that you'll need a brace of pistols and a mask.

mert - 25-8-2006 at 01:45

ok i understand....a lot thanks..there arent coins in my country

this is very bad

So,does silver nitrate produce from silver bullion in the world? is the method using in the world present?

not_important - 25-8-2006 at 06:47

Bulk silver or silver reclaimed from film processing (declining as digital cameras take over) and silver plating installations. Anyone who uses much silver reclaims it from their waste stream. Bullion is the common way to transport silver, the nitrate and acetate are the other big distribution form SFAIK, but will cost more in terms of money per gram of silver.

About the only way to get it for less than spot market prices is to expend much effort checking out second hand stores, estate sales, and similar sources, where you will occasionally find a bit of bulk silver - jewelry, tableware, decorative stuff - that is underpriced. But you need to be rather good it identifying something as bulk silver and not silver plate to make this route pay off.

If you live in a silver producing region, http://www.silverinstitute.org/supply/production.php, you just might be able to find a bit of ore that isn't being mined. This usually will be a stray vein in some forsaken wilderness, any large deposits will be obviously owned. However this is a better source of a few samples of silver minerals (read 'a few grams') than silver for production of anything.

The best way to get silver cheaply is to use your time machine; hopping back a few years to trade currency of the day for some 5 to 10 years older, until you get to about 1960.

agent_entropy - 29-8-2006 at 08:45

I was trying to think of ways to produce silver nitrate while conserving nitric acid, and this is the idea I had...

Starting with some fine (.999) silver measure out a stoichiometric excess of ammonium nitrate heat the two together in a crucible. Ammonium nitrate should melt at 170C and boil/decompose at 210C yielding ammonia and nitric acid/nitric oxides (this is still well below silver's melting point of 961C). I'm thinking that since pure ammonia boils at -33C it will be driven off by the heat leaving behind essentially nitric acid which is then free to attack the silver, thus producing silver nitrate. The silver nitrate will probably be molten as the reaction proceeds (AgNO3: mp = 212C) but should not be destroyed if the heat is controlled well enough (AgNO3 boils/decomposes at 444C). So I'm thinking, one could keep the heat on until no more ammonia is being released but stop before nitric oxides are produced by the decomposing AgNO3.

Thoughts? Improvements?

One problem I see is that it might not be only AgNO3 that is produced but also AgNO2, due to the varying nitric oxides. However, this should not be a huge problem if the overall goal is merely to get silver oxidized and into solution.

I didn't have any silver at the time so I tried just heating some ammonium nitrate to decomposition, and tested for acid by dipping some magnesium ribbon in the melt. Sure enough, the magnesium dissolved freeing lots of visible white fumes. Heating was continued until no more fumes were emitted, at which point there was white crystalline material in the crucible as well as what appeared to be red fuming nitric acid (brownish red liquid emitting fumes of the same color). Light heating was continued until the red/brown liquid was driven off and the white crystalline material remanied. The remaining white crystals were heated strongly and emitted the same reddish brown fumes (nitric oxides) which is consistent with magnesium nitrate (decomposes at 330C). The same procedure was followed using aluminum instead of magnesium, however, the aluminum did not appear to be attacked at all by the melt (a puzzling observation).

I just wanted to hear your thoughts before I try this with expensive silver.

not_important - 29-8-2006 at 16:19

Ammonium salts have been used in place of their parent acids, mineralogy field kits have done so to avoid having problems with leaking containers of acid.

There is some problem with ammonium nitrate, in that nitric acid is an oxidiser and will oxidise some of the ammonia as well as the silver; remember it's a standard method of making N2O. Given that and the potential of a runaway reaction should the temperature go much above 240 C, I'd suggest slowly adding the ammonium nitrate to the hot container+metal until the metal has dissolved.

As silver nitrite decomposes around 140 to 150 C, it shouldn't be a problem. You will want to be careful in keeping dust out of the mix, as it both reduces silver and can drag in other metals.

Aluminium is fairly resistent to HNO3.

You can always recover the silver, using zinc or the NaOH and sugar method garage chemist talked about. Heating silver nitrate, oxide, or carbonate to a nice red glow will leave silver metal.

[Edited on 30-8-2006 by not_important]

agent_entropy - 29-8-2006 at 17:13

Well, I tried the above method with a small chunk of silver metal (~0.65g) and it didn't seem to work very well, if at all. It did however leave a fairly thick coating of tightly adhered tannish crud over the surface of the silver.

Seems that the nitric acid oxidizes ammonia more readily than silver. Either that or I heated too fast or too long. I'll probably try again tomorrow; being more careful about the heating.

What about adding some HCl to the mixture? Some silver chloride might form initially, but it might help to keep the nitric acid from oxidizing the ammonia, by forming NH4Cl; which of course decomposes readily (only to re-form from the vapor phase as a fine dust), and thus would be easily driven off of the mixture.

12AX7 - 29-8-2006 at 18:37

Might be able to do something like this...

Ag2O + 2NH4Cl = [Ag(NH3)2]Cl, plus some acid and oxidizer (H2O2?) to get it going.

Once you get the Ag in solution, you should be able to fumble it around. I don't know, not sure you can crystallize something like silver nitrate without breaking the complex, causing chloride to precipitate it straightaway.

Tim

not_important - 29-8-2006 at 21:37

I think you will end up with an AgCl film on the metal. The mixed chloride/nitrate works a bit like aqua regia, great for gold and platinum but silver will form the chloride. It might be complexed and in solution, but as soon as you dilute the solution the AgCl will drop out.

The fact that you got crud on the silver makes me wonder how pure the ammonium nitrate is; fertilizer grade compounds aren't famous for being pure. You might try soaking that silver in strong ammonia solution and see what the crude does, pour the aqueous part into vinegar and see if any AgCl drops out.

I know you can recover metallic silver from the chloride by fusing it with sodium carbonate. I don't know if you can covert AgCl to the carbonate by boiling it with Na2CO3, or by bringing AgCl into solution with ammonia the adding ammonium carbonate. If that did work, then ammonium chloride + nitrate to get AgCl, convert to carbonate, boil with aqueous NH4NO3. (on its own AgCl is less soluble that Ag2CO3)

------

A varient on the hydroxide+sugar reaction : AgNO3, NH4NO3, NaOH, honey

http://dwb.unl.edu/Chemistry/MicroScale/MScale42.html

agent_entropy - 9-10-2006 at 17:58

I tried the NaOH and sugar method for silver recovery, the one that garage chemist suggested, but I can't figure out how the NaOH reacts with the AgCl. I did the gibbs free energy calculations for the reaction (net ionic):
AgCl + OH- -> Ag2O + H2O + Cl-

and it comes out to ~ +23 so it should not be spontaneous at 25C; the delta G doesn't even turn negative until about 205C (which of course is ridiculous for an aqueous solution).

I'm thinking, maybe that isn't the reaction that is happening. Any idea what is going on here?

not_important - 9-10-2006 at 21:27

Quote:

Originally posted by agent_entropy
I tried the NaOH and sugar method for silver recovery, the one that garage chemist suggested, but I can't figure out how the NaOH reacts with the AgCl. I did the gibbs free energy calculations for the reaction (net ionic):
AgCl + OH- -> Ag2O + H2O + Cl-

and it comes out to ~ +23 so it should not be spontaneous at 25C; the delta G doesn't even turn negative until about 205C (which of course is ridiculous for an aqueous solution).

I'm thinking, maybe that isn't the reaction that is happening. Any idea what is going on here?

Likely it is this

2 AgCl + 2 OH(-) <=> Ag2O + H2O + 2 Cl(-)
which is much to the left but there will be a bit on the right

Ag2O + O=C(R)-H => O=C(R)-OH + 2 Ag(s)

where O=C(R)-H is an aldehyde group in the sugar.

Metallic silver is less soluble that AgCl, and the reduction reaction doesn't go backwards very well, acid reduced to aldehyde, so the 2nd step drives things to completion.

In fact, the first step might be better as AgCl + OH(-) <=> Cl(-) + {AgOH}

guy - 9-10-2006 at 23:26

Quote:

That reaction isn't even balanced so of course youre going to be off. And make sure when you calculate the values for the anions, make sure its not the values for the GAS phase.

agent_entropy - 10-10-2006 at 04:56

Sorry, I typed the reaction in wrong, forgot the coefficients, but I did do the calculations with the coefficients (and they were for the correct phases). This is what I meant:
2AgCl (s) + 2OH- (aq) -> Ag2O (s) + H2O (l) + 2Cl- (aq)

But not important says its probably an equilibrium... hmm...
2 AgCl + 2 OH(-) <=> Ag2O + H2O + 2 Cl(-)

Also, I looked into AgOH and I can't find any data on it (even in the Merck Index) and online I keep getting redirected to Ag2O. So I'm still looking.

12AX7 - 10-10-2006 at 12:03

AgOH is a transient product. Like CuOH, it isn't stable, at least at room temperature, so it disproportionates to M2O. Probably fair to use the energy of Ag2O instead of AgOH (which would be followed by the energy of 2AgOH = Ag2O + H2O anyway).

Tim

guy - 10-10-2006 at 12:14

Quote:

Originally posted by agent_entropy
Sorry, I typed the reaction in wrong, forgot the coefficients, but I did do the calculations with the coefficients (and they were for the correct phases). This is what I meant:
2AgCl (s) + 2OH- (aq) -> Ag2O (s) + H2O (l) + 2Cl- (aq)

But not important says its probably an equilibrium... hmm...
2 AgCl + 2 OH(-) <=> Ag2O + H2O + 2 Cl(-)

Also, I looked into AgOH and I can't find any data on it (even in the Merck Index) and online I keep getting redirected to Ag2O. So I'm still looking.

Ah Ok. I see why. Ag2O is about 200 times more soluble than AgCl so it is non spontaneous at standard conditions (1 M NaOH).

You got to use G = Go - RT ln Q

when you use different concentrations than 1 M NaOH.

~~~~~~~~~~~~~~~~~~~

Also for the silver mirror reaction (sugar and Ag2O), the reaction is not 2AgCl + 2OH- ---> Ag2O + H2O + Cl-

Ammonia first complexes with Ag+ to make it soluble. Then the hydroxide is just needed to react with the aldehyde.

[Edited on 10/10/2006 by guy]

agent_entropy - 10-10-2006 at 16:31

@ guy: But I didn't use any ammonia.

My procedure is detailed below:

5 grams of AgNO3 were dissolved in 150ml of distilled water, this solution served as simulated silver containing wastewater solution. Since normally the amount of silver in a given sample of wastewater is not known, the experiment was continued as though the amount of silver in solution was not known.

Sodium chloride was added to the solution until it appeared that no further AgCl would precipitate, thus approximating an excess addition of chloride.
Ag+(aq) + Cl- (aq) → AgCl (s)

The solution was acidified with approximately 0.5 ml concentrated HCl, and was heated to boiling until the AgCl precipitate had aggregated and the supernatant solution was clear. Another few drops of HCl were added to the supernatant solution to check for complete precipitation of AgCl, no further precipitate formed upon the addition of HCl, thus it was concluded that the precipitation was complete. Precipitation of silver from the solution as AgCl is not necessary but this step was included since silver waste often already contains insoluble salts of silver.

Heating was continued and NaOH was added slowly and carefully as each addition is accompanied by violent bubbling, even when the solution is not boiling. As the NaOH was added the solution turned red, then brown and finally black as the white silver chloride precipitate dissolved and was replaced by a small amount of fine, dusty, black precipitate which did not settle out quickly. When all of the silver chloride had disappeared further addition of NaOH produced no bubbling and it was concluded that a sufficient excess of NaOH had been added.

Heating was continued and sucrose was added slowly (in case of violent boiling due to the sudden introduction of nucleation sites). As the sucrose was added the solution turned opaque milky-green, as a gray-black precipitate formed (metallic silver). As further sucrose was added the solution turned bright orange-red, and when no further precipitate formed it was concluded that sufficient sucrose had been added.

The supernatant solution was decanted and another 150ml of distilled water was added to the precipitate and the mixture was boiled in order to aid in the removal of the supernatant solution that had been included in the precipitate. The precipitate was filtered and rinsed with several portions of distilled water and methanol in order to allow quicker drying.

The theoretical yield for the initial 5 grams of silver nitrate is 3.24g of silver. The actual mass of silver recovered was 3.20g ; a recovery of 99%.

*The two things I'm still trying to figure out are:

- the AgCl and NaOH reaction

- why sucrose was able to reduce Ag+ to metallic silver (since sucrose is not a reducing sugar), I was thinking it got cleaved into glucose and fructose but that hydrolysis reaction is described as being done under acidic conditions in my organic chem book