Improving urea's stability in cream

Greetings,
I was wondering if anyone of you does have an idea regarding the improve of urea's stability in creams O/W which may contain an other active ingredient which makes the pH acid (around 3) of the cream.
The quantity of urea is 10%. Please note urea has low stability in the presence of water, high temperature and acid pH.

Thank you in advance

Comments

  • johnbjohnb Member
    A lactate buffer can help stabilise urea in solution but it is difficult to make an informed recommendation without knowing the identity of this other active ingredient and the other components of the cream.
  • Can you raise Ph to 4 as urea is most stable at 4--8.What is lowering the PH?
  • em88em88 Member
    Hello,
    Thank you for your replies. The other ingredient is salicylic acid.
    So with the current formulation, I adjusted the pH from around 3.5 to 5.5-6 with triethanolamine (2%). What I've noticed recently is that the pH after a few days will increase to 8-8.2. This worries me about the urea's stability. 
  • johnbjohnb Member
    What is the purpose of this cream?
    If it is as a keratolytic then it would most likely be better not to neutralise the acid with TEA.

    I worked on a urea/salicylic/lactic application for wart treatment and had no problems with stability. It contained 15% or so of propylene glycol and a fatty alcohol thickener.
  • em88em88 Member
    Indeed its purpose is to be keratolytic. 
    First formulations were without TEA, but after a few hours the cream had a "cheese" appearance due to CO2 release. 
    Did the cream you prepared remain at the same consistency? Also why did you add propylene glycol?
    Thank you
  • johnbjohnb Member
    The composition I worked on was more of a translucent gel rather than a cream. The propylene glycol was present to act as a solvent for the salicylic acid and as a humectant to prevent the gel drying out before it had chance to act. It also acted as a stabiliser for the whole product. There were no problems with consistency, possibly because it was not an emulsion in the commonly accepted meaning. My formulation also contained lactic acid as a salicylic acid stabiliser and as an auxiliary keratolyte.

    I don't understand why your composition should generate carbon dioxide. What are the ingredients?

    I don't understand either why you are neutralising the salicylic acid. It is the free acid that has the keratolytic properties, not its salts. From what we have been told, I think the whole concept needs reassessing.
  • em88em88 Member
    edited May 2017
    Your formulations seems to be far more better than mine, would you mind giving a few more details? Also what was the percentage of lactic acid and salicylic acid?

    I've read several articles claiming that urea decomposes with releasing CO2, especially in presence of water, acid pH (less than 4) and at higher temperatures. So I tried to avid all these incompatibilities. I agree with you regarding the concept of the formulation, that is why I am asking for help.
    Here is the current formulation
    Cetearyl Alcohol
    Petrolatum
    Paraffinum Liquidum
    Polysorbate 80
    Sorbitan oleate
    Cetyl palmitate
    Isopropyl myristate
    Triethanolamine
    Ethilic alcool
    Glycerin - 20%
    Salicylic acid
    Urea
    Purified water - 25%

    Thank you
  • em88em88 Member
    One more question, what was the pH of the gel?
  • johnbjohnb Member
    edited May 2017
    I worked on this more than 20 years ago and so any information is from memory and may be somewhat hazy.
    The product contained:

    Behenyl aclohol                             5
    Propylene glycol                          15
    Behenamidopropyl dimethylamine 4
    Lactic acid                                     6
    Salicylic acid                                10
    Urea                                              3
    Water to                                    100

    There were a few more ingredient but not any that woul make any difference to the overall product.

    I have no idea of the pH as any reading would be meaningless in the presence of so many interfering ions so it was not examined.
  • em88em88 Member
    edited May 2017
    The formulation you have worked was based more on the keratolytic effect, urea at 3% doesn't do anything extraordinary to the skin.:)
    I haven't considered a formulation for gel containing these two active ingredients, I would assume, if I get the same results of stability, the gel would be bubbly. In any case, I will try a formulation and see where it goes.  Thanks for the kind and help!

    Have a look at this patent. https://www.google.com/patents/WO2006102004A2?cl=en They do provide some formulations using NaOH. Also I have seen some other creams (last one was on amazon) containing TEA.
    Urea is such a pain to work with. I've noticed after unsealing the cream (from one of my earlier formulations) stored in tubes, the cream will get outside by itself proving an increase of pressure inside the tube which is caused most likely by CO2. 
  • johnbjohnb Member
    You are quite right that urea at 3% doesn't do anything. I doubt that higher levels do anything real, either. The reason that urea was included in my formulation was that my client insisted that it should be present. The final product did not contain urea and the only reason I gave the urea formulation here is because you requested that I do so.

    If you can get a suitable product without urea, I strongly encourage you to do so. If the main reason for the inclusion of urea is as a moisturiser rather than as a keratolyte then there are plenty of other materials which will fulfil that need (glycerine and propylene glycol even).

    In the link you have given, the NaOH containing examples are not of any particular importance, i.e they do not figure anywhere in the Claims.
  • em88em88 Member
    Unfortunately I can't remove urea and it has to be at 10% level. This ingredient is well known as "good" active ingredient for skin conditioning by patients that's why is highly requested. While salicylic acid should be at higher level (as you suggested at 10%), but for concentrations of over 2% the cream can only be registered as drug and not as cosmetic cream. 
    Was your salicylic acid gel transparent at the end or it got a whitish color? How did you manage to dissolve salicylic acid? 
    I'm sorry am asking you so many questions, but I'm kind of new in semisolid formulation, I've worked only with solid drugs.
  • johnbjohnb Member
    edited May 2017
    The gel was translucent (as stated previously) and contained microcrystals of salicylic acid. The acid was dissolved in the glycol (with heating) and the solution dumped quickly into cold water. There is more to it than stated there but the description gives an outline of the process.

    I am becoming a little confused. Are you formulating a cosmetic moisturiser or a fully functional keratolytic?

    The product I worked on was intended for the removal of warts and may not be applicable to what you seemingly want now.

  • em88em88 Member
    Hyperkeratosis dermatitis, follicular hyperkeratosis, ichthyosis, psoriasis, acne etc I'll need both the keratolytic and emollient effect. I know that the gel you prepared is for warts removal, I asked because I may have to make a gel formulation with 10% salicylic acid too and was curious how did you manage to dissolve the salicylic acid since 1g needs around 500g water to dissolve also you did not have alcohol in the formulation which would dissolve the acid much easier. I noticed now from my tests that the increase of pH was due to TEA acting very slow. I will try to find the ration of TEA and acid to maintain the pH within 4.5-8 were urea is stable.
  • johnbjohnb Member
    but for concentrations of over 2% the cream can only be registered as drug and not as cosmetic cream.
    The conditions you identify:
    Hyperkeratosis dermatitis, follicular hyperkeratosis, ichthyosis, psoriasis, acne etc

    are beyond the definition and capability of a cosmetic and any effective product should be classified as a drug.

  • chemicalmattchemicalmatt Member, Professional Chemist
    Two things here, friends: definitely apply the lactate buffer to his at pH 5.0 - plus, and add a lot of propylene glycol, not only to solvate SalAcid but to lower the water activity stabilizing urea even more.  Also: urea at 5.0% or more softens keratin (that would be skin, right?) remarkably well.  That is why it has been used in skin-softening creams for over 80 years. Urea is one the oldest known skin-care treatments in our industry.
  • Developed product similar to this with caveat it was anhydrous.Big seller. no sting from AHA.Consumers loved the product.
  • em88em88 Member
    johnb, I know, I have just mentioned where is the cosmetic going to be pointed. Obvious it will be described as a cosmetic. Thank you!
    chemicalmatt, I can add SalAcid up to 2% and reduce TEA to try to atchive pH 5 (I guess), can you please be more specific regarding the lactate buffer?
    I have added glycerin to lower the water, do you think propylene glycol would do better than glycerin?
    Unfortunately I'm not sure I can lower the urea to 5%. Thank you!
    DRBOB@VERDIENT.BIZ, did your product have urea? Thank you!
  • em88em88 Member
    Is it reasonable to think that starting from a formulation with 1% SA and 1.25% TEA, the pH increased from 5.5 to 7.9-8 and remained at this level, that by making the following change, 1.5% SA and 1% TEA the pH should be around 6-7? Or can I go straight to 2% SA and 1% TEA and still have a pH over 5? 
  • johnbjohnb Member
    Neutralised salicylic acid does not act as a keratolytic so there is little point in using TEA (or any other base) to do this.

    TEA salicylate (trolamine salicylate) is used as a topical analgesic.
  • em88em88 Member
    Hello,
    I have to at last to increase the pH to at last 4.5...
    I found an other cream in the local market. They use urea 15%, lactic acid (not specified the % and salicylic acid as preservative (most likely 0.5%). None of the ingredient mentioned on the box should increase the pH, still the pH is 7.4
    Petrolatum, aqua, urea (15%), glycerin, plant extract, cetyl alcohol, plant etract, lanolin, glyceryl stearate, plant extract, paraffinum liquidum, lactic acid, salicylic acid. The ingredient are written from the highest % to the lowest. 
    Any thought?
  • johnbjohnb Member
    Doesn't this:
    I found an other cream in the local market. They use urea 15%, lactic acid (not specified the % and salicylic acid as preservative (most likely 0.5%).
    take us back to my first reply in this thread:

    A lactate buffer can help stabilise urea in solution
  • em88em88 Member
    Yes, I've red that and was going to try (it's a great idea), but currently I don't have LA to make the buffer and it will take some time till I receive it, can I use an other buffer? For example phosphate?
  • em88em88 Member
    Ignore my last question. I did a research and found out that the lactic buffer isn't very suitable, while the citric buffer has better ranges. 
  • johnbjohnb Member
    I believe it is the lactate ion that is more important here, not the pH value.
  • @em88 we did not use urea nor did we use any base 
  • This post is apt!

    I am not making any keratolytic product. No need for low pH.

    I like to know how strong or what the capacity a buffer is required to stabilise a given amount of Urea? Is there a calculation for it? Let's say I want to use:

    a. Lactic-Lactate buffer

    b. Citric-Citrate buffer

    Is temperature factored in? I understood that Urea decomposes quicker as temperature rises.

    Simplified calculations are appreciated. 

    I read some articles including patents, they do not mention about how many grammes of buffer or acid is used. As if pH is the only thing that matters.

    If pH is the only thing that matters then I guess 1% buffer would do, as long as pH is between 4 and 8 as @DRBOB@VERDIENT.BIZ  mentioned. 

    In case you may wonder what my other ingredients are that may affect the total/final pH, please focus on only Water, Urea, and buffer. 

    If you insist, my 'secret' is:

    Urea 20.000 %
    Propanediol 20.000 %
    PQ-7 (~9.5% active) 2.000 %
    Water q.s

  • PharmaPharma Member, Pharmacist
    edited May 21
    @johnb Allegedly, it's the pH but honestly, lactate buffer has its upper limit at pH ~5. Hence, choosing lactic acid-lactate at pH 5-6 to stabilise a solution from getting alkaline sounds completely stupid (nevertheless, it's used in different Eucerin products... LoL).
    @Cst4Ms4Tmps4 Temperature is always sort of an issue with organic buffers because they change pH with temperature ;) .
    As a rule of thumbs:
    A: Add enough to capture all the extra acid/base you expect to appear.
    B: If you need more than 10 % of the compound you want to buffer, try other options. Correctly, it's % molal but being a rough estimate, % weight does just fine.
    C: There's always a formula to calculate stuff. The rule is simple: 2 lactic acid molecules capture 1 degraded urea, 2 citric acid molecules capture 3 degraded ureas. Though, you don't need to capture 100% of the urea but just trace amounts which would otherwise rise pH and cause exponential urea degradation/pH-rise. If you need a lot, than your urea preparation is so unstable that you should reconsider your formulation.
    D: Every buffer has it's range. If you expect a pH rise, you should opt for one which has it's pKa at or less than 1 unit (+0,5 would be perfect) above the desired pH of your preparation. This way, you have between 50 and 90% of the added buffer really acting as buffer and get the most out of it without actual pH changes.

    @Theotherusers Read something about a company which sprays aqueous urea solution on corn starch and let it dry. Urea crystals form within the starch and can then safely be used in their creams.
  • PharmaPharma Member, Pharmacist
    Darn, it made a smiley out of my D doublepoint :)
  • @Pharma Oh dear! It is 6.02 x 10^23?  :s

    How do I calculate "2 molecules" of something capturing "1 or 3 molecules" of another thing without going through the insane equation? I believe there must be simplified version of it. Perhaps give some examples or tell me the name of it so that I can, hopefully, find online calculator for it.

    The calculation for neutralisation does not seem to work. Maybe I do not know how to 'twist' it about. 
  • PharmaPharma Member, Pharmacist
    @Cst4Ms4Tmps4 By using molecular weights.
    But don't forget: this is cosmetics, not exact science and there's no real use for such a calculation ;) . If you want to mix a buffer, go with a buffer table and use a pH meter (that's also what scientists do first). If you want to add a preventive buffering ("quenching" would be a better word) system, you may want to go with something that shows additional effects and simply add "enough" if them. Here's a little list of such "double-edged" ingredients:
    - Lactic and pyroglutamic acid and their salts are humectants
    - Citric and phytic acid and their salts are sequestrants
    - Anisic, levulinic, and benzoic acid and their salts are preservatives
    Use a mixture of two, three or even more and chances are that you get a product which is buffered over a broad pH range. Simply adjust the final product to the desired pH and you're (probably/hopefully) good to go.
    That would be my advice and it would still be my advice even if you were a rocket scientist ;) .
  • @Pharma

    Why, very informative! Looking at Phytic Acid's molecular weight, I already fall in love with it! Its molecular weight is big AF! It is not something I could acquire, unfortunately.  ;)

    Sadly, there is no example of calculation.  ;)

    Do you mind showing just one calculation? Say, how much Citric-Citrate buffer to neutralise how much Urea. I guess it is Ammonia that is going to be neutralised, not Urea.  ;)

    Citric acid should be good as I need a sequester.  But like you said, use a mixture of two, three or more. The more the merrier by getting each of their different additional effects other than merely buffering. ;)

    Just one calculation. Just one example of the calculation. Pweety pwease. Your effort is certainly highly appreciated. ;)
  • PharmaPharma Member, Pharmacist
    edited May 24
    @Cst4Ms4Tmps4 Okayokay, I'll do it... but first some digression (to build up suspense and trying to give you an understandable picture of the processes involved and why that formula isn't really helpful).

    Imagine you're having a party (your product) on a flower field (your vessel). On there are people (water molecules) and bee hives (urea). You're playing music (pH) and the people have fun and are dancing around. Obviously, sooner or later, someone stumbles against a hive and the bees (ammonia) start swarming (urea gets degraded). The more bees flying around, the more people start panicking, knocking over more bee hives... and before you know it, the party's over (your product is toast).
    In order to protect people from bees or vice versa, you could physically separate them like I mentioned in an earlier post about corn starch. Unfortunately, bees like flowers (urea is super water soluble)...
    Another option is to calm the people by bringing acid to the party (lower the pH) and avoid playing heavy metal (high pH) which would have them pogo dancing and knocking over the hive even faster. Too bad that acid will render the bees quite sensitive, having them swarm out at the lightest knock against their home. The perfect balance would be pH 6.2 as for example recommended by the DAB/NRF (German guideline for pharmaceutical preparations).

    An important point in an earlier post is water activity. Sadly, it got more or less completely ignored. Water activity is as follows: Take a glass of tap water (full water activity) and pour it over your head ;) . Now, take some ice (frozen water has nearly zero water activity) and try to get wet: even an iceberg the size which could sink a Titanic wouldn't work. Surely, a cream doesn't contain ice unless it's ice cream but by adding solutes, the water molecules gain order and form sort of a liquid crystal structure. It's like inviting a few singing hippie girls (solutes such as lactate salts) to your party. People around them will gather together, holding hands and listening. These flocks around your flower power girls aren't static, people still move a bit, can change hands, join or leave the groups. The more girls there are, the less people are wandering around and stumbling against the bee hives, water activity drops and the party gets more stable.

    Still, it will never be at a point where there are no disturbed bees unless you don't invite people (avoid water completely). Therefore, you hang out some glue traps (use a buffer). As you can see, there's no need to distribute one trap for every bee. Besides, and that's probably the more important point, hives/urea does not just decompose to bees/ammonia but also to carbon dioxide, a gas... errr.... let's call it bee farts. A glue trap might catch the bees but won't prevent them from letting a last fart.

    In theory, carbon dioxide (CO2) reacts with water to form carbonic acid and that one reacts with two ammonia to form ammonium carbonate -> problem solved, everything neutralised. Too bad that only ammonia is well water soluble, CO2 on the other hand is very slow to react with water and gasses off. As a rule of thumbs, a gas has a 1'000 times greater volume than a solid. In other words, if 0.1% urea in a 5% urea solution degrades, the volume increases by 5% (5 ml in a 100 ml tube). The resulting foam or bubbles will be noticeable and increases pressure enough to turn your product unsellable. Therefore, the human body uses carboanhydratase, an enzyme which catalyses (speeds up) the conversion from CO2 to carbonic acid and vice versa. Without it, you would be dead within 1-2 minutes. Synthetic catalysts doing the same job aren't, for many reasons, suitable for cosmetic preparations. This means that there's no way of getting rid of CO2 gas. It also means that the very small quantity of ammonia which might appear before your product explodes (a fully degraded 5% urea solution would turn a 100 ml tube into a 5 litre balloon) is very easily buffered away, likely by also present salicylic acid (bar tenders) which are poorly water soluble but dissolve well in alcohol(ics). Alcoholics form similar flocks around them like the other people around the hippie girls. Still, there will be enough salicylic acid available to catch all ammonia (bar tenders are great with fly swatters). Now, why doesn't adding salicylic acid suffice? Because for one, it's an acid and hence probably too much (see above), and for another, it means that water activity is VERY important, probably more so than anything else. Adding acid or a buffer (the acid part of it) just keeps degradation from speeding up but not from happening in the first place.

    Patience, we're getting there... slowly...
    In order to be able to calculate things, you have to know that it's only a 1+1=2 on paper or if pH doesn't matter. As soon as you require a certain pH or pH range, buffers become important. Well, in creams and the like, there's often enough inherent buffering capacity to do the trick in everyday stuff. A buffer works by catching acids and bases alike without changing pH too much. This effect is only apparent within +/- 1 unit around the pKa for acids and pKb for bases. Every acid or acid moiety has a pKa (i.e. citric acid has 3) which is a constant: It means if pH equals pKa, half the acid is in acid form (protonated), half in salt form (deprotonated). Shifting pH by 1 unit changes the 1:1 ratio by a factor of 10. In case of lactic acid where the pKa is 3.8, at a pH of 4.8 1 part will be in acid form and 10 parts will be in salt form. At pH of 5.8 it's 1:100 and at pH 6.8 1:1'000. Urea is most stable at pH 6.2 or at 1 lactic acid to 250 lactate. In other words, only 0.4% of the added lactic acid buffer are available for capturing ammonia (on the other hand, the whole rest would neutralise an added acid). Theoretically, lactic acid is a useless buffer for a pH 6.2. Also, lactic acid is less acidic than salicylic acid and hence, the buffer composition of a mix of salicylic acid plus lactic acid-lactate buffer will in reality be comprised of lactic acid and salicylate (mostly just salicylate) unless a lot of non-aqueous liquid is added to shift equilibrium and pKa values (which are solvent dependent).
    If you still wonder (you still with me?) why adding lactic acid-lactate buffer does work: It's not about the buffering (capturing ammonia), it's merely the pH (and that one won't really change before the bottle explodes) and water activity (cause lactic acid and lactate are good in lowering water activity).

    Got the picture?

    (Your post is too long, more in a second one.)
  • PharmaPharma Member, Pharmacist
    Finally, here's the formula part:
    1 urea gives 2 ammonia and one citric acid captures up to 3 ammonia. We neglect pH here, else it would be a mix of 8 citric acid to 92 citrate at pH 6.2 resulting in ~50% hydrogencitrate and ~50% citrate. Simply said, max 50% of the added acid would be able to capture a single ammonium before the buffer is used up completely. Notably, a triple acid makes calculations VERY difficult, so difficult that either programs are used to predict things or people go with tables/graphs based on measured values (I did exactly that) and estimate from there.
    No offence, but for you, I go with the simple stuff because calculating is useless anyway:
    Urea has a molecular weight of 60.06 g/mol and citric acid monohydrate (to usual form) of 210.14 g/mol. "g/mol" is a measure of grams per a certain number of atoms. We could also work with molal which reflects the amount of acid/base moieties per molecule but that's confusing and doing a rule of three as follows is easier.
    We want to know how many grams of urea correspond to how many grams of citric acid: 1/60.06*210.14=3.5 (rounded)
    3 urea = 6 ammonia = 2 citric acids molecules: 3.5/3*2=2 2/3
    In a 5% urea preparation of 100g are obviously 5 g urea and therefore 11.66 g of citric acid would be required. Your starting solution would be at pH ~1.5 and at the end of full degradation around pH 7 (neutral because triammonium citrate is "self-buffering").
    Now that you know why you don't need that much buffer but rather a stable product to start with, this huge quantity of citric acid doesn't just sound stupid from a cosmetic point of view.

    I hope you enjoyed the good night story and learned something useful today.
    Greez!

  • PharmaPharma Member, Pharmacist
    edited May 24
    Oh, BTW I forgot to mention: To do the whole thing buffered at pH 6.2, one would require 145.75 grams of citric acid and some base (I am to tired now to calculate how much alkali one would have to add in addition) to still end up with a pH of 7. Obviously, that wouldn't fit into a 100 ml container, no would it?
  • @Pharma
     ;) I love the way you explain it!

    I think you might have misunderstood my intention (or question). I do not mean to suck up all the products of Urea degradation! 145.75 grams of Citric Acid is insane!

    To make the buffer with the same 145.75 grams, it will be 145.75 grams of Citric Acid monohydrate + 2247.1329 grams of Trisodium Citrate dihydrate. Oh yea, 2.2471329 kilograms. This makes 84 M buffer.  :D

    I do know what buffer is for and I do know that its usage is very little. That's why it is called buffer! However, I do not know the calculation and the deeper understanding of buffers, the mixes, the reactions and degradation, the explosiveness, and so on as you explained in frighteningly well and entertaining manner. (I darn sure keep this link for my future reference. I never let the effort of Tom, Dick, and Harry to go to waste, and then I ask the same bloody questions again and again the next time. I detest people doing that to me, thus I do my best not to do the same to others)




    My plan is to utilise maximum amount of Sodium Citrate, at 1.5% or 3% because it seems all right with cheap lame electrolyte-intolerant Carbomer. I guess it is due to its high molecular weight compared to other salts (Sodium Lactate is the worst offender in my case. This is the very thing that made me disabled for a very long time, wasting time going natural only to fall back to Carbomer...once again). I compared them and the difference I noticed is their molecular weights. The maths could be correct about this.

    At 1.5% Sodium Citrate--required Citric Acid will be 2.32%. This makes 0.13 M buffer.




    I am using chart, guide, calculator! Because I am illiterate I had to go through Internet nonsense in order to find the correct and simplified ones. There are still certain things that I need to manually calculate.
    For example is from w/v% to mol/L and from mol/L to w/v%. There is a calculator for it but is a bit troublesome that I need to click and click. Memorising simple MW of the acid and its salt and copy paste I can get the result quicker than I say Bob's a salted fish. Do it repeatedly regurgitating 12 digits is nothing, until I calculate another chemical. ;)

    Another aspect that manual calculation is still needed is most, if not all, that I notice, calculators and guides do not factor in molecular weight.




    About the formula part, I think it depends who is saying it. I see some use only ONE pKa of Citric Acid, there are also some use ALL pKa of Citric Acid to do the calculation. However, majority agree to pick the pKa nearest to the target pH that we want.

    For my case...

    1.5 % (Tri)sodium Citrate
    = 15 g/L

    15 g ÷ 294.1 MW of Trisodium Citrate = 0.051 mol/L


    With the help of this I punch in pH=6.2, pKa=6.39, CB=0.051. Required weak acid is 0.07898964756453644 mol/L. It does not mention what sort of acid it is. Looks like I need to dig it up.

    0.07898964756453644 mol/L x 210.138 MW of Citric Acid monohydrate = 16.59872655991655842872 g/L

    16.59872655991655842872 g/L ÷ 10 = 1.659872655991655842872 %

    1.66 % Citric Acid is to be added.

    Total ‭0.13 M. Should be plenty much to stabilise Urea. For how long will this capacity last...I have not the slightest idea. I hope God of maths and future tells me. No body wants exploding tubes and bottles, this is what we care after all! No wonder many people moved away from using Urea due to its particular negative characteristic. I, for one, like Urea very much because it works extremely well and it is an inexpensive witchcraft.

    By using this one as 'second opinion' and for fact-checking, both the calculations/results (or semi-automated) are nearly match made in heaven.

    Please note that the input is:

    CA c=0.07898964756453644 pKa=6.4
    SC c=0.051 pkb=7.8

    *****CA and SC do not matter. They can be anything. Do not have to be there. Only for reference. CA is Citric Acid. SC is Sodium Citrate. Only for my own reference*****

    If all pKa of Citric Acid is used (replace pKa=6.4 with pKa1=3.13 pKa2=4.76 pKa3=6.39) pH becomes 3.34.



    To further complicate things and if all 3 pKa are used, this may be appropriate. At the same 0.13 M and pH 6.2, Citric Acid (anhydrous) is 0.2256% and Sodium Citrate is 3.4777%.

    I do not know which is more 'correct'. One pKa or three pKa. What I know is three pKa requires significantly more Sodium Citrate than one pKa.



    You mentioned about water activity, this is what I realised a long time ago. Urea, usually pharmaceutical, is in lipid; ointment. As you already know this greasy/oily is highly undesirable, lots of drag can make you apply even more than needed. This can be also another reason why people generally move away from using Urea, not away from ointment per se.

    This article, page 129, states that it is preferable to prepare Urea containing preparations without buffer or acid, and use it fresh, use it within one month. This certainly is impractical and inconveniences the user and the maker. But of course in the article it mentions product amount is made just right for X amount of times and X amount of preparation to apply on which part of the body. Do you think humans are THAT consistent?  :p Not to mention there are people who want to save more money by stretching how long they can use a product, past best before date even!


  • I forgot to say that I have 20% Urea in my formulation!

    But that will change. I will reduce it to 10%. I find 20% is kind of tacky. There is nothing much on my list that is tacky.

    Urea 20.000 %
    Propanediol/Propylene Glycol 20.000 %
    PQ-7 (~9.5% active) 2.000 %
    Water q.s
    (0.5% Carbomer)

    I hope 0.13 M Citric-Citrate buffer could stabilise 10% or 20% properly. But for how long is debatable. Maybe the maths could predict something about it. Argh the maths again!

    The amount of buffer cannot be too much. 1.5% or a bit more total Sodium Citrate appears to be the sweet spot. Otherwise it is Carbomer's turn to be destabilised by the power of electrolyte.  :D What sort of juggling madness it is! I cannot win, can I!

    I know there are 'advanced' stuff out there for everything. I simply cannot spend too much on those. They either do not exist in where I live (Malaysia), they must be purchased in large quantities, or exorbitant for me (I'm not doing a business to earning back my pennies). Oh you better believe me that I was and still am exceedingly frustrated trying very hard hunting for chemicals, which is the reason why I finally settled down and utilise very simple, common, cheap, and old school, but tried and true chemicals.

    You may ask/wonder why 20% Urea (there is always somebody asking or wondering why someone uses this and why uses that). A few individuals had said that 30% is too high. I then reduced it to 20% and still too high. It is because I find that 20% works very on me.  Yes, on me and only me. I am not into profit, not yet, at least not now. No, I do not have dry crackling skin nor any skin diseases, just that high concentration of Urea seems to work well for and on me. No irritation whatsoever. Just me.

  • PharmaPharma Member, Pharmacist
    Just a short answer for now: It's basically about the pH, not the amount of buffer. Go for pH +/-6.2 (that's what we do in pharmaceutical preparations we make in the pharmacy). I don't see a perfect solution for that problem (and I'm not the only one). I know of water-free formulations containing high amounts of urea which aren't stable and science doesn't really know why.

    In my experience, mixtures of propylene glycol and urea are often tacky.
  • @Pharma

    All right. Your short answer for now is accepted. I will tell you my story.

    I think the calculated numbers are too high. No wonder it was extremely tacky!  :tongue:

    I made Propylene Glycol out of commission. I do not know why the said mixtures are often tacky. You did not mention their concentration used that makes them tacky. My educated guess is the more Hydroxyl units the more tacky it is. Glycerin (3 hydroxyl units) is loads more tacky than Propylene Glycol (2 hydroxyl units).

    I know that Urea will not work well without enough moisture that is the reason why I out in high amount of Propylene Glycol (20% is still much less tacky than 2% Glycerin if used alone. Weird).

    I also dropped Urea to 10%. Does not change the initial tackiness (Not sure why. Sort of like a ritual. Must be is the humidity). About 6 minutes and beyond, tackiness is very little to none, the longer the more undetectable it becomes. Observed timing is here, uncontrolled environment, high humidity. In an air-conditioned room it was like nothing.

    Armed with new knowledge I ought not worry not enough humectant. I learnt that salts are the most-cost effective humectants and are not tacky (if used with care) compared to polyols and possibly diols as well. No wonder why some literature state that adding tiny amount of NaCl can enhance Urea. They do not mention 'enhance' how though. I now know, I think I know the reason now.

    I also take your advice very seriously that pH is more important than buffer strength. I randomly made up a very weak salt, so weak that it gives me a very comfortable non-tacky and moist sensation, much like Sodium Lactate only without the horrid scent of spoilt milk. Because it is very weak Carbomer is happy to have it. Still liquid-ish, although high yield value still present. I dare not add more than 0.5% Carbomer lest it be tacky.

    I learnt something new during my frustrated moment. A few times I calculated to neutralise Citric Acid with Sodium Bicarbonate with good old reaction stoichiometry, that few times pH is already at about 6! No need extra step to first make Sodium Citrate and then add more Citric Acid to adjust pH!

    I was frustrated because I could not find reliable source that teaches me how to mix both acid and base in one single step making the salt and required pH, rather than doing it separately (because I have no Sodium Citrate nor Sodium Gluconate). If you could share this one-step approach, please kindly share it with me. I will like, share, and subscribe to your channel. Oh, I will definitely turn on the notification bell.  :sweat_smile:

    Henderson–Hasselbalch equation does not emphatically apply to base-acid neutralisation. It applies only to SALT and its acid/base. I do not know why some people suggest the equation. Maybe they don't know any better.


    On the other hand, Gluconic Acid is neutralised at pH 7 again and again.

    What is going on? Something to do with Citric Acid's dissociative identity disorder having three pKa? Please do not tell me again that triple acid is impossible. I know you can do it!  :tongue: No calculation required for this.

    Initially, I thought salts' pH follow their acid's pKa. Then, I thought Citric Acid has three pKa, it is only natural it is pH 6.39 and because it takes 3 molecules of Sodium Bicarbonate making Trisodium and therefore inherently chose the 3rd pKa which is 6.39. If it is Disodium then it is pH 4.76. If Monosodium then it is pH 3.13. Go ahead and laugh at my expense! :wink:

    After analysing long and hard, looking at Sodium Chloride, Sodium Gluconate, and other salts, their pH is not truly 7 but between 6.5 and 9. And then the knowledge that there is such thing as 'acidic buffer' AND 'alkaline buffer', that makes whatever I learnt about neutralisation went down the drain. Can't trust text books! This is the reason why science keeps me interested and curious! Science is about knowing, learning, and it works, what works and what doesn't work, does not require anybody to believe.


    I read and read your bees, party, girls, honey, flower field story again and again, repeatedly like a broken record in my bloody head. I sincerely and genuinely want to understand how the world works! I am not the type of half arse who gets answers, does it, goes to bed soundly, not bovvered, and case closed.

    I 'kinda' 'finally' 'get it'. Correct me if I am incorrect--

    1) Urea is a species which is sensitive to high pH. Specifically higher than pH 6.2. Specifically it is Ammonia that is sensitive to high pH. 

    2) If Ammonia detects above pH 6.2 nuclear fission happens.

    3) Below pH 6.2 nuclear fission also happens.

    4) As long as pH is maintained at ±6.2 Ammonia stays in its ivory tower and happily allows the rest to rule Camelot however their heart's delight to their heart's content.

    5) pH perfect at 6.2 still does not stop Urea from degradation. 

    6) No matter how strong or how many glue traps (buffer and buffer strength) there are, Urea is predisposed to churn out bees (Ammonia).

    I had already fully understood that all things, not limited to Urea, are decaying, nothing can stop the process. Even diamond degrades at atomic level. The reason is atoms keep moving, they rub into each other; friction is the destructive force. Same principle as to why water is still evaporating (sublimates) at freezing temperatures no matter what our lousy science text books brainwash us to believe.

    Perhaps my/our major concern with Urea is exploding containers and extremely undesirable acrid Ammonia scent. This is the sole reason why this drama is created.


    7) Only at pH 6.2 will Urea churn bees out more slowly than other levels of pH. Just slow and steady, slowly but surely.


    Wot! Science does not know why...More sleepless night and down the rabbit hole for me! LOL!


    You wrote:
    We want to know how many grams of urea correspond to how many grams of citric acid: 1/60.06*210.14=3.5 (rounded)
    3 urea = 6 ammonia = 2 citric acids molecules: 3.5/3*2=2 2/3
    In a 5% urea preparation of 100g are obviously 5 g urea and therefore 11.66 g of citric acid would be required.


    I am a bit confused there as there is no label. I am mentally blind, you see.

    - 3.5 is in gramme of Citric Acid. 1g Urea corresponds to 3.5g Citric Acid

    - 3.5g Citric Acid ÷ 3 molecules of Urea x 2 molecules of Ammonia = 2.3333333g Citric Acid

    - 5% Urea x 2.33333g Ctric Acid = 11.66g Citric Acid

    Did I correctly understand you?


    I detest mathematics. Without it I would be even more blind, ironically. I would be doing unnecessary trials, errors, and wastefulness. Before the maths I was doing by feel and by "what people/the Internet says". I was indubitably an insignificant ignorant fool.

  • PharmaPharma Member, Pharmacist
    edited May 31
    quote: Your short answer for now is accepted.
    So, here's the longer one:

    quote: To make the buffer with the same 145.75 grams, it will be 145.75 grams of Citric Acid monohydrate + 2247.1329 grams of Trisodium Citrate dihydrate. Oh yea, 2.2471329 kilograms. This makes 84 M buffer.
    No, not like that. Only 145 g citric acid and then adjust pH with some alkali. I was too lazy to calculate everything ;) .

    quote: I guess it is due to its high molecular weight compared to other salts (Sodium Lactate is the worst offender in my case... I compared them and the difference I noticed is their molecular weights.
    And the fact that citric acid is a triple-acid. In most cases though, several charges on a molecule make things worse (e.g. bridging molecules and have them gel or precipitate). Lactic acid is fairly small and hence "hard" whereas citric acid has a larger surface and is hence rather "soft". Each molecule has preferences for "hard" or "soft" partners (now I have to think of flowers and bees... I wonder why  o:) ).

    quote: For example is from w/v% to mol/L and from mol/L to w/v%. There is a calculator for it but is a bit troublesome...
    That's why I created some excel files with all the stuff I have in stock and the needed formula.

    quote: I see some use only ONE pKa of Citric Acid, there are also some use ALL pKa of Citric Acid to do the calculation. However, majority agree to pick the pKa nearest to the target pH that we want.
    Depends on what you're trying to achieve. If you're simply trying to keep pH stable at about one of the pKa's, than that one pKa is all you need because the other acid-parts will not budge (stay nearly fully protonated or deprotonated).
    Buffer tables and calculations are only meant as estimate! pH depends and is affected by many things such as temperature, salt concentration... In the end, one tries to add just not enough of the acid or base (in your case NaOH), measureing pH and then slowly-slowly adding the remaining brine to bring pH up exactly to where it needs to be.

    quote: It does not mention what sort of acid it is.
    A buffer is as good as always composed of a strong acid and a weak base or vice versa. Mixing strong with strong will result in a salt; that's why NaCl (table salt) has more or less no direct effect on pH. A mixture of weak/weak like ammonium acetate will result in a salt which can dissociate ammonia and acetic acid. It's also not a good buffer because you have two "flexible" things. It's better to have a "static" one as a constant/fixed point. If you take citric acid (weak acid) and add X parts of NaOH (strong base), X parts of the acid will be in salt form. If you added ammonia (weak base), it would be less than X parts because not just citric acid is buffering but ammonia too and it's all wobbly-wobbly.

    quote: Do you think humans are THAT consistent?
    I had a customer at the pharmacy for whom I regularly mixed an OTC eye cream with additional 10% urea.

  • PharmaPharma Member, Pharmacist
    Regarding tackiness:
    Quats (the standards are choline chloride and glycine betaine) and some organic acids such as citric and lactic acid tend to show strong melting point depression when mixed with urea or polyols. If the ratio is optimal, they form so called deep eutectic solvents which are usually of syrupy consistency. Take for example glucose and fructose, dissolved alone in water, they're simply dissolved sugars but when mixed 1:1, they form artificial honey which does not dry out on your skin but remains tacky. The advantage is, that compounds can only penetrate skin when in liquid form and given that an eutectic mixture doesn't dry out, it can fully absorb. The drawback is, that such a mixture remains tacky till it did whereas a "normal" dissolved salt will simply dry out and crystallize on the skin and therefore doesn't feel tacky.
    Multi-component mixtures are tricky to predict whether or not they show such an effect and it may appear only during the process of drying/resorption because composition changes when certain ingredients penetrate faster than others. NMF is such a mixture as well as nectar and honeydew... every plant knows how to concoct such deep eutectics whilst scientists just re-discovered them recently and still struggle do create mixtures of more than 3 components. BTW deep eutectis with urea render urea probably even more prone to hydrolysis...
  • @Pharma

    A quick one.

    This only gets more and more interesting.

    Is Triethanolamine Citrate a good buffer (neutralised Citric Acid by Triethanolamine) ? Is the Triethanolamine part a weak base and is another buffer on the alkaline side making the buffer wobbly-wobbly?

    My understanding of "salt" and "ions" is suddenly lost since I discovered that Triethanolamine does not 'split' into smaller components in neutralisation as NaOH or NaHCO3 does.

    NaOH --> Na+ + OH-

    NaHCO3 --> Na+ + HCO3-

    Triethanolamine --> Triethanolamine 

    And since Carbomer is sensitive to metal ions, I thought that Triethanolamine the most suited to neutralising Carbomer and be a buffer in Carbomer.

    In my test though, Carbomer turns to liquid all the same, does not matter whether Triethanolamine Citrate or Sodium Citrate. LOL! However, there is a slight improvement with Triethanolamine Citrate. Very slightly. Few people said why bother if it is only very slightly. In response to that I said we use Carbomer for a few valid positive reasons.

    I guess that having ONE sodium based base to neutralise Carbomer and ONE Triethanolamine based salt in buffer is better than having TWO sodium in buffer and to neutralise Carbomer. Sodium ions accumulate. I can be wrong.

    Sad, no chelator to chelate sodium and monovalent metal ions. Maybe there is, maybe is polyamine, but I doubt it is worth getting it.
  • PharmaPharma Member, Pharmacist
    No, it's not a good buffer. It's not the worst depending on the pH range you wanted to buffer. For a "cosmetic" pH, one would consider using either citric acid + sodium/potassium citrate or triethanolamine + triethanolamine HCl or the like (mineral acid for neutralisation).
    Sodium, carbon dioxide and water do not split either, their reaction product NaHCO3 does. Salt is the combination of an acid and a base; reacting TEA with an acid such as hydrochloric acid results, once dissolved in water, in TEAH+ + Cl-. If you simply add TEA to water, it still acts as base and abstracts a proton from water: TEAH+ + OH-. It's like ammonia. The one you buy is dissolved in water and actually consists of NH3 + H2O <-> NH4+ + OH-. The stronger the base, the more the equilibrium shifts to the right. Ditching Na2O (sodium oxide) into water doesn't give an equilibrium Na2O + H2O -> 2 NaOH and that's what we call a strong base, no or nearly no equilibrium.

    TEA can neutralise and hence gel carbomer, if used as base. But that's mixing medium strength acid with medium/semi-strong base where both, acid and base stand in equilibrium with their "free" form. Using NaOH or KOH results in carbomer salts which are good buffers, a polymeric and hence weird and hard to predict one but nonetheless good buffers. Carbomer triethanolamine on the other hand is not as neat. It works and it may suffice for cosmetic intentions.
    Carbomer (depending which one) may stay/turn liquid if not enough or too much base (depending which one) has been added or it falls out. Hence, neutralising carbomer with soda lye and adding a buffer has carbomer "feel" like using too much neutralising lye. See, using the exact amount needed neutralises at best 100% of the acids on carbomer, applying more lye simply adds "salt" but doesn't neutralise more because there is no more to neutralise.  Adding more lye or additional salt simply interferes with electrostatics, the whole trick of carbomer is annihilated. Carbomer depends on a certain amount of negative charges (salt form) to gel, if these charges are shielded by floating around positive charges from added salts no matter that there are also negative ones floating around, the effect is gone, carbomer molecules collapse like they weren't neutralised, and the product liquefies.

    Citrate carries 3 negative charges and might thereby bridge carbomer salt molecules. This effect is weaker than the aforementioned shielding because it's yet another ion layer further away from the carbomer molecule (carb-...TEAH+...citrate3-...TEAH+...carb-).

    Besides, for carbomer it doesn't matter if there is 2 Na+ or 1 Na+ and 1 TEAH+, all that counts are total positive charges and these are for both cases 2.

    Gotta go... nighty-night!
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