Sciencemadness Discussion Board

Silver Nitrate Calculation/maths check

CHRIS25 - 10-10-2012 at 05:16

Hi, yes it's me again and when I have sufficient confidence in my own rusty brain I will stop asking people to check. Please could you have a look at this:

Ag+2HNO3=Ag(NO3)+NO2+H2O

107.8+126=169.87+46+18

But for 30% nitric acid:

357mls instead of 126mls.

I have 11.4g of silver to waste. so therefore:

11.4/107.8 x 100 = 10.6%. 10.6% of 357mls = 37.8mls

So for 11.4g of silver I should use 37.8mls of 30% HNO3. The theoretical yield being 17g of Ag(NO3). But I have a feeling this last bit is wrong, I simply did 10.6% of 169.87.

vmelkon - 11-10-2012 at 08:45

If you start with 11.4 g of Ag.
Convert to moles
11.4 g / 107.8 g/mol = 0.1085343 mol of Ag

You will need 2 times the amount in HNO3
0.1085343 mol * 2 = 0.21706864 mol of HNO3

Convert the above to grams
0.21706864 mol * 63 g/mol = 13.67532432 g of HNO3

If 30% HNO3 means 30 g HNO3 per 100 g.
13.67532432 g / 30 g * 100 g = 45.5844144 g of your 30% HNO3 is needed.

It depends on what 30% HNO3 means. Is it in terms of weight or volume?

CHRIS25 - 11-10-2012 at 08:55

Hi, well I have no idea because the Bottle simply says 30%, no w/v or vol or weight words anywhere. Thankyou for the above I'm going to have to go through that again. While you are here? Do you think you could tell me if I have this right: Wanting to make 5% citric acid from the solid. I actually took 5% of 72g because 72g is the solubility in 100ml water meaning that that is the concentration. So i added 3.6g per 100ml water. Is this the correct procedure? I am not needing gram for gram accuracy, but acceptable accuracy say +/- 5%, which I know is huge in tiny quantities, but I am dealing in half litres most of the time.. (since this is all to do with photography). Thankyou.

kavu - 11-10-2012 at 09:24

A 5 % mass-volume solution (the common one) contains 5 g of material per every 100 ml of solvent.

vmelkon - 11-10-2012 at 09:32

There is a large variety of concentration expressions
http://en.wikipedia.org/wiki/Concentration

I was assuming your HNO3 30% was using mass fraction.

CHRIS25 - 11-10-2012 at 11:30

Quote: Originally posted by kavu  
A 5 % mass-volume solution (the common one) contains 5 g of material per every 100 ml of solvent.
Oh! But what if a concentrated solution is say 4.5g in 100ml? Then surely 5g is certainly not 5% strength. 5% acidity means 5% concentrate and each acid has a diifferent weight of dry matter (gas/solid) to make it concentrated, hence logic, at least my logic suggests that the initial amount of dry matter will vary considerably for each acid, different weights but all 5%?

vmelkon: "There is a large variety of concentration expressions" Thankyou for that pointer.

[Edited on 11-10-2012 by CHRIS25]

[Edited on 11-10-2012 by CHRIS25]

[Edited on 11-10-2012 by CHRIS25]

kavu - 12-10-2012 at 05:22

Quote: Originally posted by CHRIS25  
Oh! But what if a concentrated solution is say 4.5g in 100ml? Then surely 5g is certainly not 5% strength. 5% acidity means 5% concentrate and each acid has a diifferent weight of dry matter (gas/solid) to make it concentrated, hence logic, at least my logic suggests that the initial amount of dry matter will vary considerably for each acid, different weights but all 5%?


Then you can't make a 5% solution and that's it. This is why molarity is the best way of presenting concentrations for stoichiometric use.

arsphenamine - 12-10-2012 at 08:09

Given the abundance of nitric acid density tables indexed by %weight and temperature, I'm betting that a %HNO3 concentration is a weight-per-weight figure.

This isn't an analytical lab prep, so most values get only 3 or 4 sig figs.

30% %HNO3 (aq) has a density(20C) of 1.18 g/ml.

1.18 g/ml = 1180 g/litre
30% of it is HNO3 by weight, or 354 grams/litre

HNO3 Formula Weight is 63 g/mol (neglecting isotopes).
30% = (354/63) moles/litre = ~5.62 Molar or 5.62 M

11.4g Ag is 11.4/107.9 or 0.105 moles.

In a simple world in which ions at all solution strengths have a 1.0 α activity coefficient, you'd need (2 * 0.105) = 0.21 moles of HNO3 to dissolve that silver.

moles / (moles/litre) = litres, so 0.21 / 5.62 is 0.03737 litres or ~37.4 ml.

37.4 ml 30% HNO3 is the bare minimum needed to dissolve 11.4 g silver.
At a 1.18 g/ml density, that volume weighs 44.1 grams.

CHRIS25 - 12-10-2012 at 16:13

Thankyou I see that I at least had the original quantities correct though by a different means. However I will have to go through your maths and that of vmelkon's and see why you both differ. At least I learn.

arsphenamine - 12-10-2012 at 18:28

Quote: Originally posted by CHRIS25  
Thankyou I see that I at least had the original quantities correct though by a different means. However I will have to go through your maths and that of vmelkon's and see why you both differ. At least I learn.
vmelkon made a few small transcription errors.
My calculations are the One, the Truth, the Infallible (note capitals).

Look on my works, Ye Mighty, and wonder, "...like, what the fuck?"

More seriously, the crucial bits are how 30% is a weight-per-weight spec and the HNO3 density is 1.18 g/ml or 1.18 kg/litre. It all flows cleanly after that.

-hizself

CHRIS25 - 13-10-2012 at 04:21

Phenamine - thankyou, "More seriously, the crucial bits are how 30% is a weight-per-weight spec and the HNO3 density is 1.18 g/ml or 1.18 kg/litre. It all flows cleanly after that." good to learn this.

ElizabethGreene - 15-10-2012 at 16:26

Hello.

With the disclaimer that I am an utter noob, isn't the correct procedure to titrate the nitric acid? This gives you a solution with a known moles/liter value? Or do I not understand the function of titration?

-elizabeth

ElizabethGreene - 15-10-2012 at 16:33

Related subject question: When the silver is dissolving in the nitric acid solution, it liberates NO and NO2. Could those gases be drawn through water to reduce the hazardous waste and reclaim some of the Nitric Acid?

tetrahedron - 15-10-2012 at 17:12

sure, you would recover some nitric acid that way, but only the NO2 hydrolyzes producing nitric acid (and more NO), so make sure you add enough air to convert all NO to NO2, and find a way to contain the NO.

Confirming Silver or Silver nitrate method for production

CHRIS25 - 20-10-2012 at 02:42

I synthesized the silver alloy to a solid piece of 6 gm worth of silver as a lump. To check that it was free from copper I re-dissolved it in 20ml of 30% HNO3. The solution is perfectly clear, no perceptible blue at all, so job well done. I now have a choice, I either boil the whole solution to leave me with solid silver again, or allow the solution to slowly evaporate to leave me with Silver Nitrate crystals. Is this correct? The boiling away idea because as I have read on this site so often, that the nitrate decomposes quite easily at high temperatures above the 70c mark, plus the water will obviously disappear.

[Edited on 20-10-2012 by CHRIS25]

[Edited on 20-10-2012 by CHRIS25]

kristofvagyok - 20-10-2012 at 14:10

Quote: Originally posted by CHRIS25  
I have read on this site so often, that the nitrate decomposes quite easily at high temperatures above the 70c mark, plus the water will obviously disappear.

Silver nitrate is quite stable, it is also meltable (mp. above 200 according to my memories), it will decompose on lower temperature in presence of organic impurities.

from 11g to 6g to 4g.....So glad I didn't start with Gold

CHRIS25 - 21-10-2012 at 01:27

I startted off with 11.7g of silver alloy, after the first synthesis I ended up with 6.1g. After re-dissolving and evaporating I have 4.1g of silver nitrate. This is not good chemistry:(

Allowing for the fact that my silver might well have been 725 and not 925, (there were only hallmarks, no numbers) that would mean that I started with 8.5g of silver. So I lost 2.4 along the way. Then I lost another 4 during evaporation??? I'm lost. Can someone help please before I wreck my biggest chunks of silver.

vmelkon - 21-10-2012 at 03:39

No, you don't lose anything during evaporation. You probably weighted your silver nitrate and it was humid and then you weight it again, it is going to be lighter.

I'm just guessing here. You don't give much details about what you did.

CHRIS25 - 21-10-2012 at 06:07

Basically, 11.4g silver alloy dissolved in 37mls HNO3. Via copper displacement I then filtered the residue of silver mud with distilled water 3 times and collected the mud. While wet I blow torched until there was only white residue left, all the brown mud dissappeared. I then blow torched this whitish mud on a pre-charcoaled piece of wood with a much higher temperature at 1200c. I obtained a nice piec of silver at around 6.1g. I then dissolved this in HNO3 looking for traces of copper impurities. The liquid was clear. I then allowed this to cool and evaporate naturally and weighed it while it was just a little hydrated, not fully dry. It is only 4g. This is a brief overview, I think something is wrong along the way, I feel I should not have lost 2g from the silver solid to the silver nitrate. If I lost anything at the first stage, when it originally went into the HNO3, before seperation, then I know probably when this happened, I think at the filtration stage, maybe lost a gram or two there, I know why. But coming from 6g solid silver to 4 gram silver nitrate seems odd?

I think I may have one explanation: When it first cooled there was still the solution, I picked it up and after a few seconds of shaking the beaker the crystals formed very quickly, quite a lot and the liquid seemed to dissappear. After leaving this for half a day, half of the crystals dissappeared and the solution re-appeared. I then did something that I think I should not have, I assumed, without thinking, that this was all there was and poured the solution away to leave only the precipitated crystals, which after about a few hours seemed to settle and diminish a little?? This last observation is purely subjective. I was wondering whether or not he silver nitrate would re-dissolve partly? And that is where I lost my 2g??

[Edited on 21-10-2012 by CHRIS25]

vmelkon - 21-10-2012 at 15:06

If you go from 11.4 g (silver alloy) to 6.1 g (pure silver) then that means 6.1/11.4 = 53% which means you still had silver nitrate in your solution.
The problem with displacement reactions is that when you place you stick of metal in there, it gets covered with your second metal and the reaction slows down a lot (or stops).

I don't have experience with your case but from what I have seen, you can shake the copper and the silver crystals fall off down to the bottom of the solution and new crystals begin to grow.

kristofvagyok - 21-10-2012 at 15:28

So, let's see an old method from book written in 1954 named "pharmaceutics and their preparation".

The silver is granulated by the following: they melt it and pour it in large amount of water with stirring (this is not so easy, so please don't try it at home). These granulates (15kg) is placed in a large bath containing 20liter of 50% nitric acid. After the reaction happens, under circa 6 hour (beware of foaming out), they turn up the heating and boil it for a while.
When nothing happens, the upper layer is poured to another flask and the remaining silver is washed with large amount of water.
The discarded solution what contains the silver is boiled till dryness and heated further. The AgNO3 will melt and the Cu-nitrate will decompose. This will happen fully at 250 Celsius. The resulting blackish stuff is added to a lot water, stirred until no more dissolves and the remaining CuO is decanted, the water clear solution is evaporated to dryness to get the remaining AgNO3.

Good luck.

cyanureeves - 21-10-2012 at 15:29

add salt to the left over silver nitrate solution and get your silver chloride then reduce it with muriatic acid and zinc penny cores.

CHRIS25 - 22-10-2012 at 00:17

OK ok ok guys stop. First, the silver chloride method sucks. I have read about it and it's an unnecessary step. Secondly, vmelkon and kristof I was past the first stage, I know where I might have lost the first part, in the initial seperation as previously explained. I am more concerned to know whether or how or why or what with regards to the loss of two grams from the dissolving of the solid PURE silver to the now dry silver nitrate that is sitting on my desk.

And Kristof:======== When nothing happens, the upper layer is poured to another flask and the remaining silver is washed with large amount of water. The discarded solution what contains the silver is boiled till dryness and heated further. ====== This is very bad English from 1953, it is ambiguous. The words "discarded" and "upper layer" refer to the same liquid? it seems here that both the discarded solution (or is this the the upper layer?) and the upper layer contain silver which makes no sense. As I said this is confusingly written.

kristofvagyok - 22-10-2012 at 02:49

Quote: Originally posted by CHRIS25  
The words "discarded" and "upper layer" refer to the same liquid? it seems here that both the discarded solution (or is this the the upper layer?) and the upper layer contain silver which makes no sense. As I said this is confusingly written.


Yes, they should have meant the same.

CHRIS25 - 22-10-2012 at 04:44

Oh well. Anyway, I'm going to make some more silver nitrate and then buy an ounze coin. It really works out cheaper than buying silver nitrate from a supplier. I stained my fingers! I only touched the sides of the porcelain dish when I was fiddling - lesson learned.

Displacement of silver by copper in nitric acid

CHRIS25 - 24-10-2012 at 15:02

Hi, Since I lost about 4 grams of silver the first time around I do not want to make a mistake with my larger piece of silver. Hence this question: At the moment I have plenty of silver precipitated out of solution via copper. Now the reaction is still ongoing very very slightly with plenty of what now is quite clearly a very less dense precipitate (what is precipitating now has the same bouyancy as the solution, it tends to float rather than sink straight away like the precipitate of the first few hours did); Question. Is it possible to still have silver in the copper/silver nitrate solution that will not be displaced by the copper because the copper does not have sufficient nitrate left to dissolve and kick the silver out for want of better words?

Hexavalent - 25-10-2012 at 08:59

Theoretically, the reaction would continue going until you ran out of either the silver nitrate or the copper. You can speed up the reaction somewhat by gently heating, and adding more copper to your mixture.

CHRIS25 - 25-10-2012 at 13:57

Quote: Originally posted by Hexavalent  
Theoretically, the reaction would continue going until you ran out of either the silver nitrate or the copper. You can speed up the reaction somewhat by gently heating, and adding more copper to your mixture.


Thanks Hexavalent, I just wanted to be sure. I have a dark green solution of copper nitrate, PH2, and a blue solution, PH3, of copper nitrate from the previous silver precipitation experiment, am I right in assuming that the greener one is far more saturated with Cu2+ complex than the blue one? I am just curious.

Also can I titrate in order to discover how much copper there is.

And I read somewhere ages ago that I can use Ammonia to discover whether there is any silver left in solution. Now I know this will be minimal since I ran both experiments to their completion, IE overnight, I added two drops of ammonia into a test sample from each solution and immediately saw a bright deep blue just before dissolving. I also diluted a test sample and then added drops of ammonia from a syringe; the first few drops produced a milky reaction before the next few drops which then turned to that bright deep blue. I would like to understand more of what is happening, it's not essential for my photography, but I want to learn since this can only benefit my appreciation for investigating off-shoots and links to various reactions. thankyou.

kadriver - 7-11-2012 at 04:46

I know this is an old thread, but I could not resist adding some input since I am new to this forum.

CHRI$25,

It looks like you may have had excess nitric acid from the dissolution of the silver - this will cause the silver mud (it is actually called cement silver because it looks like cement) to float as you described earlier in this post.

What is happening is as the silver "cements" out from adding the clean copper to the silver nitrate solution, the excess nitric acid begins to immediately redissolve the cement silver. globuals of cement silver get full of gas as it dissolves, causing the globuals to float.

I have encountered this problem and this is what I did to solve it.

Use just enough nitric acid per gram of silver to dissolve it. This will ensure that most or all of the nitric acid will be consumed during the reaction and your cement silver will have a bright crystaline appearence. Plus you save nitric acid for use with other experiments.

When you dissolve silver metal with dilute nitric acid, first weigh the silver and then use that weight to calculate the amount of nitric acid you will need to completely dissolve the silver.

It takes about 1ml concentrated nitric acid to dissolve 1 gram 925/sterling silver.

It takes about .7ml concentrated nitric acid to dissolve 1 gram 999 fine silver.

To dissolve, first put on a pair of nitrile or latex gloves, silver nitrate stains skin on cantact for 3 or 4 days. Then put the silver in a beaker and cover with a watch glass.

Using the numbers from about, calculate the amount of nitric acid needed to completely dissolve the silver - keeping in mind that the amount needed will be different for pure silver (999 fine) and 925/sterling silver.

Then use this calculated number, and add that amount of DISTILLED water to the beaker with the silver in it. Don't use tap water because it frequently containes dissolved chlorides (chlorine) and this will combine with the silver nitrate to form silver chloride in your reaction vessel causing loss of your metal.

After adding the calculated amount of distilled water to the silver, replace the watch glass and place the container on low heat.

Measure out the calculated amount of nitric acid and begin adding a few ml of nitric acid to the container with the silver and distilled water. Add it slowly and incrementally using a pipette. As the reaction slows, add a little more acid until the silver is completely dissolve. I like to have a tiny piece or two of silver metal left in the bottom of the reaction vessel. This acts as an indicator to let me know that all the nitric acid has been consumed.

The key here is to let the reaction take its time. Fresh acid will make the reaction move along much quicker, but then it is much easier to over shoot the acid - the very problem we are trying to avoid.

If there is a few undigested pieces of silver left, then what you can do is pour off the silver nitrate solution, and add drops of nitric acid to the remining bits to dissolve them completely.

A small amount of free nitric acid is required in order to get the silver to cement out onto the copper. If the nitric is completely consumed, then the copper will not react and it will just sit there with no silver forming on it. If this happens, then add a few drops of concentrated nitric acid, stir with a glass rod, and re-introduce the copper.

925 and sterling silver, when dissolved in dilute nitric acid, will produce a blue to greenish solution due to the copper and nickle.

999 fine silver should produce a clear colorless to faint pale yellow colored solution with no traces of blue or green.

When cemeting silver with copper, you can tell when the reaction is complete when the silver no longer forms on the copper. As long as silver continues to form on the copper, there is still silver in solution. When complete, the copper will appear bright pink after all the cemented silver has been shook off of it. Unless there are platinum group metals present, these metals will also cement out, usually at the very end of the cementation process. They will appear as black powder adhereing to the copper and sometimes cling stubbornly to the copper.

Another test to see if all the silver has come out of the solution is this: draw a sample of the solution though to contain silver and add a few drops of hydrochloric acid or a few grains of table salt to it. If silver is present, then a cloud of silver chloride will form very quickly. It the solution stays crystal clear, then it is safe to conclude that all the silver has been removed.

Hope this helps anyone looking to work with silver.

kadriver - edited once to add weight information

[Edited on 7-11-2012 by kadriver]

CHRIS25 - 7-11-2012 at 10:06

Kadriver - This is very very helpful indeed. thankyou for taking the time to explain something. At least now I understand the floating situation. I extracted 20g of silver second time round and only lost about 3 to 5 grams this time. The second time worked very well without any floatations this time and I learned the HCL test a few days later after the first test, and it was clear of silver. So all in all I was able then to get my silver nitrate. But your explanation has 'cemented' a few lose ends, your technique I will try next time round. But the first time yes indeed I had too much NO3, not due to calculations, but due to the fact that I mis-calculated the amount of alloy copper in the sterling silver.

Purifying Silver Nitrate

CHRIS25 - 10-1-2013 at 05:36

Had a spot of bother re-crystallizing silver nitrate and was wondering now if I had done it correctly, I found only one old reference to this on the web. My method was simply to re-dissolve in distilled water and evaporate through very gentle heating, but my result is now different from the remaining silver nitrate that I have from the original reaction. Namely the crystals are white and very tiny, colour is white and texture a bit like sugar, whereas the silver nitrate crystals from the initial reaction that I did not re-dissolve are translucent and the crystals are bigger/chunkier. I tried to find references that would help me at least to know what I should expect to see from re-crystallizing but could not.

woelen - 10-1-2013 at 06:19

I think that your fine crystals are quite good. I have commercial AgNO3 and this also is white.

The original product consists of larger crystals and usually these crystals look more like glass or ice. The smaller the crystals, the more white the solid looks like.

MrHomeScientist - 10-1-2013 at 07:29

Really the difference is just due to growing conditions. If you let your solution evaporate on its own, without heating, over a much longer period of time, you'd get larger crystals that would be more in line with your commercial sample. I've done this myself. Woelen is right in that your smaller crystals would look more white simply because they are smaller, and have more surfaces to reflect light.

CHRIS25 - 11-1-2013 at 07:27

Both of you thankyou. I appreciate that you took the time to help. I was a bit concerned that I might have contaminated somehow - a spot of rain fell for a few minutes even though the reaction was 75% covered outside, managed to cover it after a minute or so but was worried. Silver nitrate being precious could not afford to throw away and start again.

So, unable to find info on this topic, I need to ask why re-dissolving and evaporating purifies? What would the chemistry behind this be? Why would the crystal structure change like this? Sorry I can not work this one out by myself. Maybe if you have any references I could read then that would be appreciated.

[Edited on 11-1-2013 by CHRIS25]

watson.fawkes - 12-1-2013 at 07:52

Quote: Originally posted by CHRIS25  
So, unable to find info on this topic, I need to ask why re-dissolving and evaporating purifies? What would the chemistry behind this be? Why would the crystal structure change like this?
Recrystallization (a term for which you can search) is a purification process because the compound being purified has more affinity for the crystal structure than impurities do, that is, if you're lucky. Some impurities co-crystallize; for these, recrystallization doesn't work.

When you say "crystal structure", you have a terminology problem. Crystal "structure" refers to the microscopic, geometric arrangement of the molecules in a lattice. The macroscopic shape of the crystals formed is referred to as the "crystal habit" (another term to look up). Habit is an interesting topic in it own right. The most obvious change of habit is that larger crystals, as a rule, are the result of slower evaporation. Impurities in the solution can change the habit, as well.

AJKOER - 12-1-2013 at 08:32

As many silver salts are sensitive/reactive upon exposure to sunlight, and also to CO2 and sulfides as present in air, exercise caution. Your purification attempts may actually be counter productive.

CHRIS25 - 14-1-2013 at 03:06

Thanks watson fawkes. I can always count on you to highlight the nitty gritty details of things that would never occur to me. Irritating as it can feel, I actually appreciate your little illucidations and have learned a great deal from you.

AJkoer, thankyou, that is something I had considered - but from the little I read seemed for photographic purposes that this was an absolute must - the silver nitrate needs to be purified, now I know that I will never get absolute purity, Bakers analysis style if you are familiar with that, but having synthesized this silver nitrate from its silver copper alloy, (two serviette rings), I knew that the initial product would certainly not be pure enough for the old photography that I do (1840-1860). Since it will be mixed with ferric oxalate (which I still have to make) and tartaric acid (which I have made) in such small amounts I am hoping that it will not degrade its effectiveness in noticeable ways.