Chemical Forums
Chemistry Forums for Students => Analytical Chemistry Forum => Topic started by: ivychen1989 on May 02, 2006, 04:20:30 AM
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For school, I am supposed to design an experiment. I've chosen to do a quantitative analysis of Ca2+ in Milk, comparing the concentration of Ca2+ after the milk has been heated to different temperatures. Does anyone know if there would be any problems in me letting the milk cool before doing titrations? (I don't want to crack the glassware if I change the temperature in the conical flask too fast when titrating.)
Please say if someone knows of anything terribly wrong with the back titration I've come up with:
1. 10.0mL of milk and 20.0 mL of standardized EDTA (about 0.05molL-1) mixed thoroughly together so that all the Ca2+ complexes with the EDTA
2. 20mL of KOH added to bring pH up to 13, so that the Mg2+ in the milk precipitates as Mg(OH)2 (and will not interfere with the titration)
3. 2mL of hydroxyl naphthol blue indicator (turning the solution blue)
4. titrate with standard calcium chloride solution (about 0.05molL-1) until the colour of the solution changes to violet (when the extra Ca2+ begins to form a red complex with the indicator)
Would I be able to keep the KOH (for a few weeks) or will it react with the CO2 in the air? (I'll be able to keep the other solutions,right?)
Would the temperatures 20, 40, 60, 80, 100 degrees Celsius be reasonable to heat my milk to?
If anyone knows of any scientific work related to this topic, could they post it up?
Thanks for any help/suggestions.
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While there is nothing particularly wrong with the titration I wonder what you will be really determining in the experiment. Amount of calcium per sample will not change, that's mass conservation law. Perhaps during heating some of the calcium can get binded to peptides present in the milk, but you will most likely destroy this equilibrium adding strong base.
Thus it will be very surprising for me if the results will show any dependence between concentration and heating - but keep us posted, I will be happy to be mistaken on that one ;)
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this has also been done before, u can take a look at others' attempts
http://www.chemicalforums.com/index.php?topic=7158.0
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I've only been allowed to do a rough trial using beakers and measuring cylinders, so I haven't been able to tell if there are any differences in the calcium concentration. The method works, as in I've been able to figure out a concentration similar to the one written on the milk bottle).
I'll write more when I've done my proper, accurate experiment.
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i was wondering around how much calcium chloride solution in mL's is needed???
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i was wondering around how much calcium chloride solution in mL's is needed???
Depends on amount of calcium in the milk, concentration of EDTA and concetration of CaCl2.
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Jus need a rough estimate really but it's the same concentrations and volumes as is in the procedure at top of page (practically repeating the experiment)
thx
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Well, right now I've only used one bottle of milk, and have yet to try it all out on another. I also titrated after heating milk at 30, 50, 70, and 90 degrees Celsius to see if there was a trend in my results (the calcium concentration decreased then increased again as heating temperatures increased). My results are as follows:
20 degrees 0.0311 mol/L
30 degrees 0.0310 mol/L
40 degrees 0.0305 mol/L
50 degrees 0.0297 mol/L
60 degrees 0.0289 mol/L
70 degrees 0.0291 mol/L
80 degrees 0.0294 mol/L
90 degrees 0.0302 mol/L
100 degrees 0.0308 mol/L
Does anyone know/can guess reasons why this would be?
I'll put up my results for the next bottle when I've finished.
i was wondering around how much calcium chloride solution in mL's is needed???
As for how much calcium chloride solution is needed, definitely no more than 15mL for each titration, and I made 1000 mL for the whole thing, though I may run out during my analysis using the second bottle, because I did extra titrations for 4 different temperatures (and I had not taken that into account when I decided on the standard calcium volume I needed).
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Today (Saturday) I went to school and finished off my investigation. Here are my results for my second bottle of milk. The trend shown in both bottles are similar, although the calcium concentrations are higher from my second bottle.
20 degrees 0.0313
30 degrees 0.0311
40 degrees 0.0305
50 degrees 0.0297
60 degrees 0.0293
70 degrees 0.0296
80 degrees 0.0299
90 degrees 0.0305
100 degrees 0.0313
I really can't think of reasons why the trend is like this.
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What was the complete procedure? And I mean not only just determination, but whole procedure starting from the moment you've bought the milk ;)
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Well, I first made all the solutions I needed, the EDTA and calcium chloride. I standardized the EDTA I made with a magnesium sulfate solution that I'd made, and I made my standard calcium chloride solution by dissolving calcium carbonate in hydrochloric acid.
I bought my first bottle of milk at the dairy, and I took it home and put it in the fridge. Then I carried it to school the day after, and put it in the school's fridge about an hour and a half afterwards (it's been so cold lately and the classes didn't have the heaters on, so I don't think this would have affected anything).
Throughout that week, I titrated the milk. Depending on how many titrations I had time to do (one, two, or three different temperatures), I would pour out the milk that I needed into a large beaker; the rest of the milk stayed in the fridge.
To heat the milk, I poured some milk into another beaker and put the beaker on a hotplate (which was far from my large beaker of milk). Once a thermometer showed that the milk had reached the temperature I wanted, I would take the beaker off the hotplate and placed it in a cold water bath.
Once the milk's temperature had decreased to 20 degrees Celsius, I would pipette the milk into conical flasks that already had the EDTA, KOH and indicator in them. I titrated four times for each temperature to firstly get a rough idea of the titre values, then the other tree titrations gave me concordant results.
I bought my second bottle of milk on Thursday, and kept it in my fridge until this morning, when I brought the bottle to school, where I kept most of it in the fridge and only took out enough milk for me to heat at three different temperatures. So I went back to the fridge to fill up my beaker twice after that.
I did the temperatures in this order both times: 20, 40, 60, 80, 100, 30, 50, 70, 90 degrees Celsius.
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So basically milk was stored for some unknown time in undefined, albeit low temperatures, than it was heated for few minutes to the higher temperature, cooled and titrated?
I am surprised you have found any dependence. The only thing I can think of is some additional equilibrium taking place between Ca2+ and other substances present in the milk, equilibrium that is somehow affected by the heating (and protein denaturation?).
Note that amount of calcium in your samples was always the same, that's mass conservation law. Only thing that can change is free calciium concentration, or, more precisely - concentration of calcium that can be titrated with EDTA at the pH, temperature and speed of your titration procedure.
Interesting.
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One could infer that there is something in milk (protein? or fat?) that sequesters calcium when heated to a specific temperature. How often did you replicate each temperature treatment? Since the temperatures observed are not naturally occurring in animals one wonders what the importance of this analysis considering how pasteurization is done.
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So basically milk was stored for some unknown time in undefined, albeit low temperatures, than it was heated for few minutes to the higher temperature, cooled and titrated?
I am surprised you have found any dependence.
Yeah, but I made sure that I did all the titrations before the milk's best-before-date. And other students in my class have been doing the same type of investigation (though using different brands of milk); out of the three I spoke to, two have gotten similar trends to mine, and one has gotten no trend.
How often did you replicate each temperature treatment? Since the temperatures observed are not naturally occurring in animals one wonders what the importance of this analysis considering how pasteurization is done.
Each temperature was done once for each bottle, so twice. I did this investigation because of the old-wives-tale that we shouldn't boil milk because heating it decreases the calcium concentration, which is, of course, bad.
Thanks for the suggestions, I'll see if I can find anything in the milk that might have made the calcium unable to react with the EDTA.
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I did this investigation because of the old-wives-tale that we shouldn't boil milk because heating it decreases the calcium concentration, which is, of course, bad.
As the amount of titratable calcium goes up in the higher temperatures, tale seems to be wrong according to your results.
But it doesn't mean your results support the idea of not heating the milk to 60 deg C where you have found minimum concentration - as you don't know what amount of calcium can be freed when the milk is digested.
I digged some and this may be of help:
http://www.afns.ualberta.ca/Courses/Nufs403/PDFs/chapter2.pdf
(it supports the original tale in a way - although one need to check the solubility of calcium phosphate in gastric juices).
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solubility of calcium phosphate in gastric juices
Sounds like a Ph.D. thesis to me!
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Yeah, but I made sure that I did all the titrations before the milk's best-before-date. And other students in my class have been doing the same type of investigation (though using different brands of milk); out of the three I spoke to, two have gotten similar trends to mine, and one has gotten no trend.
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Yeah, that last dude...he's the organic chemist of the bunch. ;D
So, what I've to say doesn't deal with your titrations directly, but does answer some of the tangential quetions raised. The presence of phosphate along with calcium will hinder its absorbtion, and also calcium is absorbed by the body exclusively in the stomach. My assumption about the decreased absorption in the presence of phosphate is the insolubility of calcium phosphate, but in an acidic environment, one might expect to see solubilization. The other thing I know is that the presence of citrate with calcium aids in its absorbtion (the brandname of calcium suppliment Citrical utilizes this) compared to either calcium carbonate or calcium phosphate. So, drink some orange juice with your tums as opposed to pop :P
So in retrospect, this post is overall not really contributing to the thread, but whatever. It's neat stuff, perhaps you will find it interesting, even as an aside.
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solubility of calcium phosphate in gastric juices
Sounds like a Ph.D. thesis to me!
Assuming gastric juice is mostly HCl and its pH is about 2 that's just an equlibrium question - Kso for Ca3(PO4)2 and CaHPO4 are known as well as all three dissociation constants for H3PO4. Not an easy question, but not that difficult too.
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solubility of calcium phosphate in gastric juices
Sounds like a Ph.D. thesis to me!
Assuming gastric juice is mostly HCl and its pH is about 2 that's just an equlibrium question - Kso for Ca3(PO4)2 and CaHPO4 are known as well as all three dissociation constants for H3PO4. Not an easy question, but not that difficult too.
I suppose you're right...I figured there were so many matrix effects that it couldn't be calculated. But alas, chemists and their assumptions save the day :P
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Do you know the accuracy and precision of your results? What are the errors associated with them? Could it be that all of your concentrations are the same within errors?
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Do you know the accuracy and precision of your results? What are the errors associated with them? Could it be that all of your concentrations are the same within errors?
I haven't calculated the percentage or absolute error of my result, but if the concentrations were all the same within errors, I would not have gotten a similar trend with both bottles of milk that I analysed (?). And I don't think my results were all that inaccurate, as I got three concordant titre volumes for each sample (heated to a certain temperature), and these ranged from 12.49 mL to 12.97 mL (which is quite a large range?). I tried to make everything as accurate as I could in my investigation.
Of course, if you can think of anything I might have done to make my results so inaccurate, please tell me and I'll include that in my report.
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Plot your results with the error bars to see if the trend holds or not.
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Plot your results with the error bars to see if the trend holds or not.
How do I go about finding the error? Do I start by saying, for example, for making the magnesium sulfate solution used to standardize the EDTA, 0.01g error from the electronic balance used to measure the solid, and 0.5% error (?) for the volumetric flask used? And find all the uncertainties that I can think of from the equipment, and from parallax error, etc.?
Wow, that would be so complicated!
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Have you used the same EDTA solution all the time? If so, even if your results are not correct in terms of absolute numbers, their ratios are OK. Start assuming that your accuracy of volume measurement is half of the scale. Alternatively - you did three titrations for every point, so you can estimate both average and deviation for every measurement.
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Have you used the same EDTA solution all the time? If so, even if your results are not correct in terms of absolute numbers, their ratios are OK. Start assuming that your accuracy of volume measurement is half of the scale. Alternatively - you did three titrations for every point, so you can estimate both average and deviation for every measurement.
Yes, I used the same EDTA.
Plot your results with the error bars to see if the trend holds or not.
When I said my titrations were concordant, the ranges for each point were mostly 0.04 mL. If I only use that to calculate my uncertainties then plot everything on a graph, I have a trend. But I still can't explain it. I only know that the calcium hydroxyl phosphate in the casein micelles (a milk protein) is affected by heating, but it doesn't support my trend. When heated above 75 degrees Celsius, calcium orthophosphate is made, which doesn't dissolve easily and precipitates with a pH higher than 7 (my titration was done at about pH 13).
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can sum1 post how the calcium concentration were worked out?
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can sum1 post how the calcium concentration were worked out?
This is a back titration. Figure out the concentration of the Ca2+ in the conical flask that you put in (your standard solution). Then change that into moles (you know the volume you used). From the volume of EDTA titrated, figure out how many moles of Ca2+ reacted with this (a 1:1 ratio). Subtract this from the original moles of Ca2+, and you'll come up with the moles of Ca2+ actually in the milk. Use the volume of milk you put in to the conical flask to figure out the concentration of Ca2+.