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This topic comprises 2 pages: 1 2
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Author
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Topic: Anybody Good at Chemistry?
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Randy Stankey
Film God
Posts: 6539
From: Erie, Pennsylvania
Registered: Jun 99
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posted 12-23-2018 04:39 PM
I'm now working as a lab tech in an electroplating shop.
A main part of my job is to keep the chemistry of the baths under control. (i.e. Performing tests and making additions to ensure that the solutions are up to par.)
Most of the time I perform tests according to what's laid out in our procedure manual but there are a few baths that don't have any tests laid out in the procedure manual. Basically, when the bath stops working, we dump it and make up new chemistry.
I guess that's okay but it can be expensive to throw out chemicals when you don't have to. Besides, waste treatment is a big P.I.T.A!
We have one bath made up from ammonium bifluoride and hydrogen peroxide that is used to strip tin from steel. Flushing ammonium bifluoride down the drain is a big no-no. It has to be treated first. The only reason the bath goes bad is because the hydrogen peroxide degrades over time.
I discovered it's possible to titrate for the concentration of peroxide using potassium permanganate. So, I did the math and figured up a procedure that works. Now, instead of dumping 50-plus gallons of peroxide/bifluoride solution down the drain, I can use the titration to figure out how much peroxide is needed to bring the bath back into spec. We can probably double or even triple the life of the chemistry and, thereby, reduce our waste by a third or more and save money on chemicals to boot.
Problem is, I don't have a degree in chemistry. I took chemistry in college and I did fairly well at it but that was over 30 years ago. I'm really rusty!
In order for my test to be included in the procedure manual, it has to be validated. I can get it validated if I take it to one of the professors I know from the Chemistry Department when I used to work at Mercyhurst University but I have to have everything written up, T's crossed, I's dotted and all the math has to be correct. The prof. won't sign off if it's not all correct and ready to go.
I've done the write-up with all the theory and calculation that I think is necessary. I've written down a step-by-step procedure. All I need to do is perform some validation trials with peroxide solutions of known concentrations and compare them to the expected outcome.
Before I go to the trouble and expense of making 20 titrations to validate the procedure, I'd like to have somebody look over my work.
Does anybody, here, understand this?...
5 H2O2 (aq) + 2 MnO4- (aq) + 6 H+ ==> 5 O2 (g) + 2Mn2+ (aq) + 8 H2O (l)
Would you like to look over my work before I potentially waste a lot of time and chemicals on a test that doesn't work?
Reagent grade KMnO4 solution costs about $25.00 a liter. ![[Wink]](wink.gif)
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Marcel Birgelen
Film God
Posts: 3357
From: Maastricht, Limburg, Netherlands
Registered: Feb 2012
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posted 12-24-2018 12:16 AM
I'm not going to pretend to be good at this, but once in a while, I needed this damn thing called Chemistry to solve some puzzles.
Not pretending to know anything particular about your specific solution either, I came up with this:
Let's put the hydrogen peroxide into water, which is still pretty simple. It will yield the signature H3O+ couples that make up our acidy solution. In your case, this is what you already should have in your existing solution, which you're going to pimp, but at least it explains where the "H+" (or rather H3O+) came from:
(1) H2O2 (aq) + H2O (l) => HO2 (aq) + H3O+ (aq)
Then we're going to dissolve the potassium permanganate into the water, leading to a nice pink solution. Since it's not a simple salt, but a more complex compound, I had to look this one up, but apparently, it dissolves in water like this:
(2) KMnO4 (s) + H2O (l) => K+ (aq) + MnO4- (aq) + H2O (l)
So, then we're probably getting to the business end of things, although I'm not really sure what has been achieved in the end. The next one should be a redox reaction, based on all the agressive stuff in the water. If you see actual bubbling occur, let's hope it's actual O2 which is escaping.
(3) 2 MnO4- (aq) + H3O+ (aq) => 2 O2 (g) + Mn2H (aq) + H2O (l)
I guess that this Mn2H thing is pretty unstable in water and henceforth will probably end up like:
(4) Mn2H (aq) + H2O (l) => Mn2- (aq) + H3O+ (aq)
Disclaimer: Monday Morning armchair Chemistry at work here.
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Randy Stankey
Film God
Posts: 6539
From: Erie, Pennsylvania
Registered: Jun 99
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posted 12-24-2018 09:38 AM
Just for the record, the place where I work has no chemicals that can explode on their own. You'd have to mix things together that shouldn't be mixed together. If you did something like that it would be a clear case of negligence.
Any idiot knows that it's bad to mix sulfuric acid and sodium cyanide!
The shop where I work doesn't do any chrome plating. No hexavalent stuff. We do tin, solder, nickel, copper, gold and silver. We use a lot of cyanide but we don't have anything like acetic anhydride.
Like I said, you'd have to do something pretty stupid to cause a building wide disaster and there's nothing I can think of that could cause an explosion without doing something that any reasonably intelligent person should know not to do.
To explain the reaction that I'm working on:
Yes, the hydrogen peroxide (H2O2) and the potassium permanganate (KMnO4) will both be in aqueous solution. The H2O2 will be the sample taken from the solution tank under test. The KMnO4 will be a pre-made solution bought from a chemical supplier.
Both H2O2 and KMnO4 are strong oxidizers but KMnO4 is stronger so it will actually oxidze peroxide, even though peroxide is also an oxidizer. If you look at the equation it says "5 H2O2 + 2 MnO4." This means that five units of peroxide are reacting with two units of permanganate. That's the key to the whole thing.
If you know the concentration of one of the two substances, you'll be able to calculate the concentration of the other.
Since we buy potassium permanganate solution from a supplier who certifies that it is of a given concentration (0.02 Molar) we can deduce how much peroxide is in the sample.
The Manganese in potassium permanganate can exist in several oxidation states. (Having an excess or deficit of electrons in its atoms.) Under neutral pH conditions, it stays at oxidation state VII but under acid conditions it can be reduced to state IV. When it's a IV, it becomes clear in solution.
So, if we put some sulfuric acid into the mix, we add those extra "6 H+" that you see in the equation above and this will ensure that the purple permanganate turns clear when it reacts with the peroxide.
So... If you carefully mix KMnO4 into a solution containing H2O2 and a little bit of H2SO4, it will turn clear until all of the H2O2 is used up. At that point, the solution will no longer turn clear. If you do it until your reaction flask JUST turns to a very faint pink/purple tinge, you have reached the balance point of the reaction.
Now, you note how much of your KMnO4 solution you used and you can do the math to calculate how much H2O2 was in your sample.
We already know that it takes 2 moles of KMnO4 to react with 5 moles of H2O2 so that means the two substances react at a 1:2.5 ratio... For every 1 unit of permanganate you need 2.5 units of peroxide.
From this point, it's all just high school algebra.
So, let's say that I've done my test and I know that there is 100 ml/L of peroxide in the solution. The bath is supposed to have 125 ml/L of peroxide. That means I have to add another 25 ml/L of peroxide.
The processing tank contains 150 L and I need another 25 ml/L so that means I need to add 3.75 litres to bring the bath back into spec.
So, I go out to the shop, put on my apron, gloves, boots and face mask, pour out just shy of a gallon of peroxide and add it to the tank and we're back in business.
This is the thing I have to write up and have all the math done right before I can submit it for approval.
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Tony Bandiera Jr
Film God
Posts: 3067
From: Moreland Idaho
Registered: Apr 2004
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posted 12-24-2018 12:45 PM
Based on Randy's post above this reply, I think he works out chemistry in the same way I used to in high school, which got me in a lot of trouble....to explain:
In class, with the actual glassware and chemicals at hand, I was always the first one done with the experiments where the objective was to "mix chemical A with chemical B and write down the resulting formula and concentration of the result". I never had an issue and was right pretty close to 99% of the time.
But, when it came to homework, doing the exact same thing out of the book, I struggled badly. It got to the point that I just stopped doing the homework at all. I had asked the teacher if I could sign out a few flasks and some test tubes so I could do the experiments (using water) at home. (He refused to let me.) Of course, the teacher (one of the football coaches) got pissed and called my parents in to a conference. (That created a whole 'nother problem as my dad worked graveyard shift and my mom had to take time off work.) He accused me of being lazy and threatened that, if I were on the football team I wouldn't been allowed to even suit up for the games. (Joke was on him, I was 5'9" and 120lbs at that time.)
I found out about all this when I got home that day to two very angry parents. I explained the situation which calmed them down somewhat. My mom didn't get it, but my dad understood that, like my maternal grandfather, I was the type of person who had to do things "hands on" to work out complex situations. Reading it in a book (or on a diagram) gave the background info, but to solve it I needed to get hands on. (I am still that way to this day.)
Seeing this thread reminded me of all this.
Back to the problem at hand, :
quote: Randy Stankey
We already know that it takes 2 moles of KMnO4 to react with 5 moles of H2O2 so that means the two substances react at a 1:2.5 ratio... For every 1 unit of permanganate you need 2.5 units of peroxide.
From this point, it's all just high school algebra.
So, let's say that I've done my test and I know that there is 100 ml/L of peroxide in the solution. The bath is supposed to have 125 ml/L of peroxide. That means I have to add another 25 ml/L of peroxide.
The processing tank contains 150 L and I need another 25 ml/L so that means I need to add 3.75 litres to bring the bath back into spec.
Like Randy said, seems like simple algebra would provide the answer, and his example above seems clear.
What I don't get is what other math is needed to get the approval? Are they expecting to see the full reaction on the molecular level explained in a formula? If so, to what end?
I currently run the city's water system and have to manually chlorinate the water, to a level not exceeding 4ppm. There is a very simple calculation that gets me the required amount of chlorine solution (at a given concentration) to add to our 50,000 gallon reservoir. The regulatory agencies do not expect me to provide the math of the solution for them, just to know what the ratio of a given concentrate to treated water gets me to the desired result.
quote: Courtesy of publichealthontario.ca
C1 x V1 = C2 x V2
C1 is the initial concentration of the bleach (sodium hypochlorite) solution.
V1 is the volume of the bleach to be diluted with water. This is what you are trying to calculate.
C2 is the concentration of the diluted bleach solution you are preparing.
V2 is the volume of bleach solution you are preparing.
In my case:
C1 is usually 6.3%
V1 is the variable, how much of the 6.3% solution to get my 50,000 gallon reservoir to a desired concentration (ex. 4ppm)
C2 is 4ppm
V2 is 50,000 gallons.
The end result for this example is 3.33 gallons of 6.5% sodium hypochlorite added to 49,996.68 gallons of untreated water, which is a ratio of 1:14,999
In actual use, I round the added solution down slightly based on the level of the reservoir at the time I chlorinate, and routine chlorination is half that level (2ppm.)
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Randy Stankey
Film God
Posts: 6539
From: Erie, Pennsylvania
Registered: Jun 99
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posted 12-24-2018 06:12 PM
You're right, Tony. I've already done the hands-on stuff and I know that it works. Now, I have to go back and do all the math to prove why it works.
We're trying to be an ISO-9000 shop and everything has to be done by the book. If I want to use the test that I "invented" it has to be written up. How can somebody else who comes after me perform the same test, the same way I do it and get the same, repeatable results if it's not written down.
If I had a degree in chemistry I could simply write it up and say, "Randy Stankey Ph.D." and my word would stand but, since I don't have that degree, I have to find somebody else who does. I don't want to insult somebody else who does have a degree, who I am asking for their sign-off, by presenting them a poorly written half-baked paper that looks like a Freshman's homework assignment, written at 7:30 a.m. before the 8:00 a.m. due date. I want to be able to show my boss a well written paper that I'd be proud to sign my name to.
The math that needs to be validated is how to convert the amount of titrant used in the test into a real world value that can be used to determine what the concentration of the solution under test really is.
If I needed 14.5 ml of potassium permanganate to titrate a sample solution to its equivalence point, how do you determine what the concentration of hydrogen peroxide is?
Multiply the volume of titrant used in ml. by 8.598 to get the concentration of H2O2 in ml./L. (or by 1.101 to get oz./gal.)
So... How do I know that: ml x 8.598 = ml./L. ??
I got some beakers and a burette and figured it out, just like Tony said.
But, when the auditors come to look at my procedure manual and they ask, "How do you know that?" I have to be able to show them the math.
The trouble comes in when I have to calculate how much peroxide is needed to add. (Tony will appreciate this.)
The peroxide we buy is 35% by WEIGHT but I need to know how many liters/gallons to add. In order to calculate that, I need to know the specific gravity of H2O2 (1.13) and the molecular weight of H2O2 (34.014) Then I have to convert that into molarity (moles of substance per liter) in order to work out the chemical formula. Once I have that, I have to convert that all back into liters in order to know how much perxoide solution to add.
If peroxide was labeled in terms of molarity (H2O2 35% by weight = 11.63 Molar) I could have simply figured volume x molarity and I would be done.
But, since manufacturers insist on using archaic units to label their product, I have to do all the math to convert things from one unit to another and back again.
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Randy Stankey
Film God
Posts: 6539
From: Erie, Pennsylvania
Registered: Jun 99
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posted 01-06-2019 11:37 AM
I ran some trials to test things out and, now, I've got more questions.
I made up 3 samples of stock hyrogen peroxide diluted in distilled/deionized water: 50 ml./L. - 100 ml./L. - 200 ml./L.
I used 100 ml. volumetric flasks and volumetric pipettes to make up the test samples. (These should be accurate to 0.1%.)
I titrated each sample three times then averaged the results.
I got the following results:
Sample 50 ml./L. -- Titrant = 6.10 ml. -- Result = 52.45 ml./L.
Sample 100 ml./L. -- Titrant = 12.15 ml. -- Result = 104.47 ml./L.
Sample 200 ml./L. -- Titrant = 24.00 ml. -- Result = 206.36 ml./L.
That's almost 5% error! I was really confused until I went back and looked at the Certificate of Analysis for the batch of peroxide that I had used.
The C. of A. said that the stock solution that I used was 35.8% instead of the 35.0% that I had originally done my calculations with. (The grade of peroxide that we buy can be between 35.0% and 35.8%)
If I go back and adjust my math to account for the difference, my error goes down to 1.6%. That's not too terrible. I'm working in a grungy factory, not a university research laboratory. This is probably as accurate as it is possible to get under the conditions I have to work.
My question is whether I should change my calculations to the new numbers or leave them at the theoretical numbers I started with. I'm leaning toward leaving them as-is because, as the concentration of the stock solution varies, I want the test to account for that. (i.e. If I'm supposed to have 125 ml./L. of peroxide in my working solution but the supplier sends a different batch which has a different concentration, I'll have to adjust the amount of stock solution that I use to get the working concentration I want.)
If I want this to be any more accurate, I would probably have to go out and buy a bottle of hydrogen peroxide standard and run my trials again and this time, I'd probably have to do ten trials instead of three.
Aside from the time and expense, I don't think I'd be able to get any better than what I have already done under the conditions I work. I think it's better to say it's good enough and be done.
What do you think?
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Randy Stankey
Film God
Posts: 6539
From: Erie, Pennsylvania
Registered: Jun 99
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posted 01-06-2019 02:29 PM
We looked at a test kit like that. The problem is that its range only goes up to 10 mg./L. The concentrations we use are more like 5 g./L. We'd blow that test kit right off the end of the scale.
We actually have a lab with all the stuff needed to do most lab tests we need. I run titration tests for hydrochloric acid, cyanide and other stuff on a daily basis. There should be no reason why I can't run a test for peroxide, just the same way.
We use the peroxide in a solution used to strip tin from steel. If parts need to be reworked, they often have to be stripped down to the base metal and replated with tin. We use a solution with ammonium bifluoride, peroxide and a little bit of secret sauce to strip off the tin before replating.
As you know, hydrogen peroxide degrades pretty quickly. When it does, the tripping solution goes flat and won't work. When it stops working, we either have to throw it our or spike it up with more peroxide.
We have to be careful when throwing this solution out. Ammonium bifluoride makes hydrofluoric acid when mixed with water. We can't dump HF down the drain. It has to go through a complex waste treatment process before it is released to the sewer.
Besides, I'm the guy who has to dump out 50 gallons of this stuff, clean the tank and refill it with 50 lbs. of ammonium bifluoride and 5 gallons of peroxide.
If all that is necessary is to add some more peroxide to keep this solution working then it saves a lot of work for me, a lot of trouble in waste treatment and it saves money when we don't have to buy the chemicals.
Previously, we had been dumping this stuff and replacing it. I'm the one who came up with the idea of saving it by adding more peroxide.
The idea for the test is to determine how much peroxide to add to bring it up to spec. (16 oz./gal. or 125 ml./L.)
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