Synthesis of sodium citrate

 Abstract:

The simple synthesis of sodium citrate from sodium hydrogencarbonate and citric acid, for use in making Benedict’s reagent. This post also serves as a test for a new writing layout as I move slightly closer to a scientific paper format, and also includes a test video on my new YouTube channel which, for now, will produce a video now and then to accompany a chosen post here on WordPress.

Introduction:

Sodium citrate (Na3C6H5O7) can be easily and simply prepared by the acid-base reaction of sodium hydrogencarbonate (NaHCO3) and citric acid (C6H8O7):

 

3NaHCO3 (aq) + C6H8O7 (aq) → Na3C6H5O7 (aq) + 3H2O (l) + 3CO2 (g)

 

The solution will then be heated to the boiling point of water until enough water boils off that a precipitate forms, where from the solution will then be cooled to allow the sodium citrate product to crystallise from the reaction solution. The product will then be filtered, washed, dried, and stored in a container. I am aiming to form 47 grams of sodium citrate pentahydrate.

(I originally was aiming for 40 grams, although my calculations were based off sodium citrate being a dihydrate when in fact I discovered – half way through the reaction – that I would actually be forming the pentahydrate when searching for the solubility of sodium citrate [1] so excuse the funny numbers)

I plan to use the sodium citrate for use in making Benedict’s reagent, the reagent used to test a substance for the presence of reducing sugars.

 

Risk Assessment:

Substance Hazard information
Sodium hydrogencarbonate Low Hazard

It liberate carbon dioxide on gentle heating (or with acids). Sodium hydrogencarbonate is an approved food additive, E500, and is used as baking soda. [2]

Citric acid Irritant

Causes serious eye and skin irritation and may cause respiratory irritation.
It is an approved food additive, E330.
Concentrated lemon juice may contain 2-hydroxypropane-1,2,3-tricarboxylic acid (citric acid) up to 1.7 M, this concentration being classified as an irritant. [3]

Sodium citrate Low Hazard

Sodium citrate is an approved food additive, E331. [4]

Acetone
(propanone)
Flammable/Irritant 

Highly flammable liquid & vapour.
Causes serious eye irritation; may cause drowsiness or dizziness; repeated exposure may cause skin dryness and cracking.
Its vapour may catch fire above -20°C. For a 15-minute exposure, the concentration in the atmosphere should not exceed 3620 mg m-3. The smell can be detected by most people at about 47 mg m-3, well below the level which could cause harm. [5]

Sam’s personal safety notes:

  • Gloves and goggles should always be worn when handling chemicals to protect your sensitive eyes and skin.
  • Acetone readily evaporates at room temperature so it is suggested that you work in a well ventilated area due to the potentially harmful vapour. Also open flames should be avoided due to acetone being highly flammable. Acetone can also dissolve some plastics so avoid plastic reaction vessels.
  • Citric acid can make you skin and eyes slightly sore, but as long as your eyes are washed pretty quickly afterward and your skin washed before touching other objects, it is extremely safe to work with.

Experimental:

My hotplate stirrer was set up with a 250ml glass beaker on top. 100ml of distilled water was added to the beaker, along with a magnetic stir-bar.

IMG_8468
Figure. 1

 

I turned stirring on and up to medium power, while I dumped a pre-weighed 29 grams of citric acid straight into the beaker (28.58 grams to be precise but my large scales can only measure to 0 decimal places). I left the solution to stir until all of the citric acid had dissolved.

IMG_8469
Figure. 2

 

Once all dissolved, I added small portions of sodium hydrogencarbonate by spatula from a pre-weighed container of 38 grams of the compound. Carbon dioxide gas is a byproduct of this reaction so I must add the sodium hydrogencarbonate slowly to the citric acid solution or else risking extreme fizzing which could cause the solution to bubble over and out of the container, hence the larger than needed beaker.

Just 34.27 grams were required, but, due to the lack of accuracy of the scales I was using, small portions of the 38 grams were added until fizzing stopped and pH showed neutral, therefore negating the inaccuracy of the scales and showing when all of the citric acid had been neutralised.

IMG_8471
Figure. 3

 

I ended up adding the full 38 grams of sodium hydrogencarbonate to cause the pH paper to finally produced the green colour than indicates a pH of 7 – neutral.

IMG_8472
Figure.4

 

I placed a thermometer in the solution, held in place by the clamp seen on the right in figure. 5, and I turned the heating on and up to medium in order to get the solution boiling quickly. Once achieved, just below 100ºC was maintained by lowering the heating accordingly.

IMG_8473
Figure. 5

 

At around 80ml of liquid left, I transferred the solution to a smaller 100ml beaker to allow faster boiling.

IMG_8477
Figure. 6

 

Once down to ~50ml of solution left, a white precipitate quickly formed; this is seen from Figure. 7 to Figure. 8.

IMG_8479
Figure. 7
IMG_8480
Figure. 8

 

I immediately took the beaker off the hotplate and allowed to cool until room temperature with a lid on top to prevent the entering of dust. I ended up leaving the beaker for around 2 days in total, but even a few hours should suffice.

IMG_8482
Figure. 9

 

I set up for gravity filtration (although vacuum filtration can be used if you own the equipment, mine arrived just after I completed this synthesis slightly annoyingly). I poured any liquid into the filter paper, followed by the solid crystals that had formed – these crystals were very hard and required quite some persuasion from a few sharp metal objects to break up.

The filter paper ended up being too small so I eventually swapped it out for a larger piece of filter paper.

IMG_8484
Figure. 10

 

I then washed the filter paper with 2 portions of 5ml of acetone to dry the product, although I would recommend washing first with a similarly small portion of ice-cold water. I found it difficult to find the solubility of sodium citrate in acetone online, but did find that it was only slightly soluble in ethanol [6], so I decided it should have a low enough solubility to keep any lost product to a minimum.

IMG_8486
Figure. 11

 

I then took the filter paper out and placed it on top of some tissue paper in a glass dish, and left it to dry under a lamp for a day.

IMG_8489
Figure. 12

 

The final product was very lumpy so I enlisted the help of my mortar and pestle, reducing the particle size to a more easy to work with powder.

img_8493.png
Figure. 13

Results:

I was mostly satisfied with the new consistency of the sodium citrate so I placed it into a pre-weighed glass storage jar. The mass of sodium citrate pentahydrate came out as 43 grams.

IMG_8494
Figure. 14
IMG_8495
Figure. 15

Discussion:

My theoretical yield was 47 grams, while the actual yield of sodium citrate pentahydrate was 43 grams; therefore:

(43 ÷ 47) × 100 = 91.489

My percentage yield, when rounded to the appropriate 0 decimal places, produces a value of 91%. Although:

100 − ((43 − 2(0.5)) ÷ 43) × 100 = 2.33

100 − ((29 − 2(0.5)) ÷ 29) × 100 = 3.45

2.33 + 3.45 = 5.78

My percentage yield comes with a percentage error of just under 6%, so could be anywhere just below or above 97% or 85% respectively.

Either way, it is a decent percentage yield for me as I didn’t require all that much sodium citrate and just wanted to attempt a larger scale synthesis than needed as the reactants are cheap and easy to get hold of.

My yield was most likely lower than it could have been potentially due to a combination of not all of the citric acid reacting; the switching of the filter papers, as product will be lost in the discarded filter paper; the inaccuracy of the scales; and that when crystallising the sodium citrate, I did not put the solution in the fridge/freezer to reduce the temperature to gain more product – not done as the crystals already formed were a solid block and were hard enough to break up.

This is a nice easy preparation to get me back into chemistry and making posts this summer as I am currently enjoying a fat holiday before university. My upcoming posts might be on a weird schedule as I have recently become interested in Myrmecology and keeping ants as pets (of which I might make another blog to write about, so stay tuned!)

Thanks for reading!
(Please check out my video below. It is a new idea, aiming at an audience who are less familiar with chemistry than the audience this blog is aimed at. The quality is still a work in progress though, and I hope to improve, especially with the audio, in future videos)

 

Video:

Synthesis of sodium citrate
18thTimeLucky – Amateur Experimentalism

 

References:

[1] “CRC Handbook of Chemistry and Physics, 97th Edition”. p. 4-85
(Sodium citrate pentahydrate soluble in water 92g per 100ml at 25 celcius)

[2] “Student Safety Sheets, 2nd Edition, 2018”. CLEAPSS. p. 25
(Safety information on sodium hydrogencarbonate)

[3] “Student Safety Sheets, 2nd Edition, 2018”. CLEAPSS. p. 33
(Safety information on citric acid)

[4] “Material Safety Data Sheet – Sodium Citrate”. Colonial Chemical Solutions, Inc.
(Safety information on sodium citrate)

[5] “Student Safety Sheets, 2nd Edition, 2018”. CLEAPSS. p. 61
(Safety information on acetone)

[6]
“Lotus Illustrated Dictionary of Organic Chemistry”. p. 172.

(sodium citrate slightly soluble in alcohol)

10 thoughts on “Synthesis of sodium citrate

Add yours

    1. Thank you and thank you! I did try a new technique this time, filming what I was doing using a make-shift tripod and then taking screenshots at parts of the filming afterwards where I wanted a picture, this way getting some interesting action photos; I’m glad you approve!

      Liked by 1 person

    1. Hi Sean, sorry, thanks for letting me know! I checked through my junk folder and it seemed to have hidden itself away from me there. I’m extremely interested!! I’ll send you a proper email back in a few hours, I had just planned to go hunting for ant queens today as I enjoy ant-keeping and it rained yesterday, so hopefully there will be a nice nuptial flight going!

      Like

  1. Hello Samuel,

    I hope I could shoot you with loads of questions. Because I have lots of questions! I made my own research. Yes, I did and had. I also asked other chemists at Chemist Corner dot com. Still, I am confused and frustrated.

    The answer is I am NOT a chemist. I hope you would not evade simply because I have zero knowledge in chemisty. LMAO! So far only one chemist at a university literally ignored me because I do not have the ‘level’ required to deserve his assistance. What an utter nonsense! A university professor behaving like a typical pillock! He should resign if he does not like to be an educator.

    Here’s the thing, I am trying to make Citric-Citrate buffer, with pH of 6.2. It is to stabilise Urea. Even using calculators online I am still lost.

    (a) https://en.intl.chemicalaid.com/tools/reactionstoichiometry.php?equation=C6H8O7*H2O+%2B+NaHCO3+%3D+Na3C6H5O7*2H2O+%2B+CO2+%2B+H2O

    (b) http://calistry.org/calculate/ph-buffer-Henderson-Hasselbalch

    Buffer strength does not matter. What I need is:

    1. pH 6.2
    2. Trisodium Citrate dihydrate maximum 1.5% OR 1.5% in 100g water.

    Using site (a) I punched in 1.5 for Na3C6H5O7*2H2O
    1.072g Citric Acid monohydrate (I do not have anhydrous one) and 1.285g Sodium Bicarbonate are required.

    In mole would be:

    Convert from % to g/L, and then to mol/L.

    1.072g Citric Acid monohydrate
    1.072 x 10 = 10.72
    10.72 ÷ 210.1388 (MW of C6H8O7·H2O) = 0.0510139012881010075245504399949
    = 0.051 mol/L Citric Acid monohydrate

    1.285g Sodium Bicarbonate
    1.285 x 10 = 12.85
    12.85 ÷ 84.0066 (MW of NaHCO3) = 0.15296417186268697935638390316951
    = 0.153 mol/L Sodium Bicarbonate

    Tested the solution with universal test strip (I have no pH metre). It appears to be at pH 6. I do not even need to go extra step to making buffer, it is already a (nearly or maybe spot on) pH 6.2 buffer!
    I thought all salts are pH 7 when fully neutralised. What is going on? Salts also follow pKa of their acids? That means Sodium Gluconate is pH 3.86, and Sodium Salicylate is pH 2.97? And if it is Monosodium Citrate it would be pH 3.13, and Disodium Citrate would be pH 4.76?

    And when I use site (b) to check for pH, the result is odd.

    Here is what I input:

    CA = 0.051
    CB = 0.153
    pKa = 6.39

    pH is strangely 6.867.

    Only when both CA and CB’s mole are the same will pKa = pH.

    I hope you would explain to me what is happening in detail, as detailed as possible (otherwise I fail to push the daisies in peace with many unanswered questions!). Along with some miserable but simple (but not simpler) calculations for idiots, if possible.

    Liked by 1 person

    1. Hi Christopher!
      Sorry to not reply sooner, it is exam season currently and I have been very busy! I love helping out people who are interested in chemistry – whether or not they are a chemist means nothing to me which is how it should be – but I am very rusty on the old pKa and buffers! (physical chemistry is just not my thing, I am more of an organic chemistry guy actually!)
      Shame you have not been able to receive help yet and sorry for that experience you went through with that university professor. I would be happy to try take a crack at it but it might take me a several days if not a week in my spare time to try relearn this stuff and then have a go. It depends on whether you are okay with waiting even longer? Sorry, that is the byproduct of having 2 exams within the next week or so 😦
      In the mean time though if you don’t mind me asking, what led you to want to make a buffer? I understand it is to stabilise urea as you said, but is it part of your job or do you enjoy doing some chemistry at home for fun or…? Just curious!
      Have a nice weekend!

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      1. Ah, there you are, Samuel!

        I ‘stalked’ you a bit and found out that you are actually a young person. Your age that I saw online keeps changing so I thought it could be old stuff. Now that you confirmed my notion, I wondered why you are perpetually ‘young’! 😛

        Now that is the spirit of science! I used to teach kids science, so-called “unnecessarily advanced” topics such as the Black Hole. I was in disagreement with the principle, the principle tried to dismiss me from the workplace and dismissed me she did.

        You would still attend to your studies and exams even if I said I am not okay waiting! XD

        In all seriousness, your exams are more important than my questions, for now. Mine is not urgent. I wager that you could help even more people once you start working. At the same time I think you would be bugged by loads of work that you would no longer have much time helping people outside of your work.

        Oh. Only just now that I learnt what physical chemistry is. Before this I know only inorganic and inorganic chemistry because these two are the most talked about.

        To assist you in expediting the cracking, I did further search, I found one page which sort of answers one (or the entire) of my questions–partial neutralisation. Scroll down the page, titled “Preparation of buffers. Partial neutralization of a weak acid by a strong base”.

        http://butane.chem.uiuc.edu/cyerkes/chem102aefa07/lecture_notes_102/lecture%2026-102.htm

        It is that one-step approach to making buffer with acid and base, without using the salt separately or make salt first and add more acid later.

        1. I do not know whether or not the calculation applies to me because I do not use/have ‘strong’ acid/base for the purpose. I have common and easily accessible chemicals (usually from bakery shop/baking supply).

        2. The example given, the acid has only one pKa. Citric Acid, in my case, has three. No explanation and no example for triprotic acids.

        3. What I truly need is in grammes. By grammes I mean I need a maximum amount of Trisodium Citrate (dihydrate) or Triethanolamine Citrate. For instance, 1.5g of a salt in 100g or 100% water (add H2O to 100g).

        The page does not show the ‘reverse calculation’. Something like this one http://dl.clackamas.edu/ch105-04/calculat.htm. How to calculate if various known variables are given, and how to find one another value with and from a value. To put it simply–convert this to that, convert that to this, convert whatever to something else.

        You are an organic chemistry guy, eh. Hmm…might be useful. I think you can related to (poly)Acrylic Acid and Sodium (poly)Acrylate. Polymers are within organic chemistry, yes?

        Acrylic Acid requires neutralisation to become gel. Triethanolamine is its neutraliser. Sodium Hydroxide is the most commonly use. I chose Triethanolamine because…read on!

        Once it becomes a gel, its polymeric network collapses with the presence of VERY LITTLE amount of electrolytes. It is ultra sensitive to salts especially to divalent and trivalent metal ions. Trivalent ones are the worst offenders. Hence why I thought that having Triethanolamine Citrate would be safe from all forms of metal ions.

        Another aspect that affects Sodium Acrylate is is a salt’s molecular weight, the higher it is the less it affects Sodium Acrylate. Triethanolamine plus any acid (in my case is Citric Acid and Gluconic Acid) has a massive molecular weight, basically a combination of Triethanolamine’s and the acid’s molecular weights. It is strange that this combination does not react, nothing is released. I thought acid-base reaction is always fizzy! Bang goes all our bloody text books!

        I read somewhere before this that Sodium Hydroxide makes a slightly more cloudy gel compared to one which is neutralised with Triethanolamine. It means the polymeric network is interrupted. My educated guess is that I might be correct about the metal ions are unfriendly to Sodium Acrylate. Sodium Acrylate still gels with Sodium Hydroxide probably that is the maximum metal ions it could accept; simply enough to make it gel.

        I have no NaOH to try it out to know more than my own hypotheses. However, I replaced NaOH with Sodium Bicarbonate after looking at the stoichiometry result is the same as using NaOH. I have not tried going beyond the needed amount (according to calculation) because I am tired of opening up the vacuum food bag every time I need Acrylic Acid and quickly seal it back and suck the air out manually. Adding insult to injury, Acrylic Acid is extremely hygroscopic, it solidifies/becomes sticky very very quickly in air. At least in the ambient air here (I am not in an air-conditioned room). I live in the tropics…..this is the reason why…Humidity and heat here are the enemies to many chemicals. I threw away my previous batch of Acrylic Acid because it shrunk and turned hard (I don’t know this form is still usable or not). I used about only 10g of 100g. Such a great waste! It is not cheap.

        I enquired Sumitomo Seika, the said that they “think” that I still can use shrunk and hardened product, only that the amount of moisture that it absorbed from the air may make measurement inaccurate. According to Lubrizol, the material absorbs a maximum of 8% moisture as its equilibrium. Confusing, confusing. But…maximum 8% moisture and not deliquescent all sound good to me. If I still can use it will be best! No more money down the drain! In relation to measurement, if 8% moisture is the maximum that the material absorbs, then I do not think I need to be concerned about it. If the moisture it absorbs fluctuates on certain days of different relative humidity levels then it might be a pain to measuring it.

        Thus the “maximum 1.5g” of Triethanolamine salts seem to be the sweet spot for Acrylic Acid. I, however, need to reduce it further. 1.5g is merely my guess because I have zero basic in the calculation. Equations of neutralisation, pH, and whatnot are new for me. Permutations.

        Many chemists admit that they refer to chart rather than be bothered doing the troublesome calculations. While I agree with them, unless one works in a company that wants one to only obey a fixed standard for a standardised product, a fixed a chart is not flexible and certainly obstructive. I strongly doubt that they have thousands, if not millions, of charts showing different concentrations at different pH for different buffers. Those charts I found show different pH but the concentration is fixed.

        I do my calculations with Google Spreadsheet once I get all the equations that I need. Only the beginning is exceedingly challenging and frustrating. I can set those formulae in the spreadsheet and then forget them; once done I need to only input values and results magically reveal themselves before me.

        I do not mind you asking me as to what led me to want to make a buffer. A valid and innocent question.

        I am basically ‘unemployed’ and have been unemployed for a very long time. But I make things (DIY), grow things, and do my own one-man research experimentation. Those that I successfully researched, made, and tried, I try to share and sell them as a source of income.

        With that said, history proves that I am a failure as salesperson and I have trouble selling just about nonsense for the sake of getting money. Because I cannot tolerate inaccuracies nor imprecision, I do what I am interested in, and I gravitate towards science as opposed to pseudo, junk, and popular sciences. Many people say that I am ‘irrationally scientific’ to the extent that I irritate people, and that I deliberately make my own life impossible to live, and that I am a failure due to my decision. Whaaaat!?!?!?! To that I usually say their beliefs are irrational bollocks!

        So yes, making buffer is a little this and a little that. For fun, chemistry at home, my work, my job, my career, my, me, mine, and I. I do things that others believe is a waste of time, effort, and money. I do find that no amount of knowledge is wasted. The side effect of knowledge is power. Not only that, I also find that knowledge is transferable. The only downside of knowledge is either you use it or lose it, just like any skills. Generally speaking, no one actually knows how to ride a bicycle no matter they are professional cyclists or have been cycling for million million billion of years! None of them!

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      2. I hope this goes to the bottom; after my comment before this one. I do not know this comes up first or second because my comment before this one requires approval.

        I had to quickly write in as I just now received a response to my question on another site.

        The chemist mentioned that weak base and weak acid do not make a good buffer! I think Triethanolamine + Citric Acid monohydrate is not a good buffer despite of the combined molecular weight is massive. If it is true I will then make a buffer with Sodium Bicarbonate + Citric Acid monohydrate hydrate.

        I have no NaOH as the strong base, the product of mixing Sodium Bicarbonate and Citric Acid results in Trisodium Citrate dihydrate, which is the same as mixing Sodium Hydroxide and Citric Acid. The only thing that NaOH and Citric Acid does not product is CO2.

        Although there will be sodium ions, and metal ions will interfere with the world peace of Sodium Acrylate, I guess having Sodium Citrate is the wisest choice to go for in the name of good buffer.

        I also learnt that sodium is considered a “spectator” ion that does nothing, not even an acid. Only citrate ion is the dominant one acting.

        Triethanolamine Citrate on the other hand dissociates into Triethanolamine and Citrate, both are weak and weak. It’s all wobbly-wobbly, can’t make a ‘static’ buffer, hence bad buffer. Another example given is Ammonium Acetate is a bad buffer, because Ammonia and Acetic Acid are weak acid and weak base.

        I was also understood that if I am simply trying to keep pH stable at about one of the pKa’s, then that one pKa is all I need because the other acid-parts will not budge (stay nearly fully protonated or deprotonated). Phew, one less worry! pKa 6.39 and pH 6.20 are all I care!

        Then there is something about deep eutectic solvent, may speed up Urea’s hydrolysis (buffer won’t save Urea). I do not have to worry about this happening because the Citric Acid is so low.

        Here is the link to the conversation for your perusal. The user “Cst4Ms4Tmps4” is me.

        https://chemistscorner.com/cosmeticsciencetalk/discussion/comment/37577/#Comment_37577

        Like

  2. I hope you are doing well, Samuel.

    I did more hunting while awaiting for your grand answer. I think I cracked it!

    This is the source to my question.

    https://www.researchgate.net/file.PostFileLoader.html?id=54f9c88fd11b8b402a8b4586&assetKey=AS%3A273725033779218%401442272541972

    CTRL + F and search for “Example B”. There is only one Example B. The one which says “… conjugate acid/base pair will be generated in situ (in solution)”.

    As I do not have Sodium Hydroxide nor other strong bases, I replace NaOH with NaHCO3 as the products of their neutralised forms are identical, only without CO2 for NaOH. (yeah I know that I am repeating myself).

    I input the equation to my Google Spreadsheet, most values are automated. I tweaked a little to suit my need.

    Target final volume : 100 mL
    Target pH : 6.2
    Buffer concentration : 1 M

    Base : 84.00660928 M (Sodium Bicarbonate)
    Acid : 210.1388 M (Citric Acid monohydrate)

    pKa acid : 6.39

    After all the calculations, required are:

    – 0.4758759449 mL Citric Acid monohydrate
    – 0.1867046618 mL Sodium Bicarbonate

    Because Citric Acid is a triprotic acid, 0.4758759449 Sodium Bicarbonate x 3 = 0.5601139854 mL Sodium Bicarbonate

    I checked it with the equation for “partial neutralisation” pH is exactly 6.2!

    http://butane.chem.uiuc.edu/cyerkes/chem102aefa07/lecture_notes_102/lecture%2026-102.htm

    I also checked it with this page.

    With that said, I presume the number might be incorrect, although the calculations are correct. It is because it is in mL. I always use g because I never first dilute anything and mix later, if at all. How do I convert all mL to grammes? Something to do with Molality? I suddenly am brain dead looking at the mass of numbers not knowing where to start changing the values into grammes! XD

    One more thing, is the aforementioned equation does not allow me to pick the amount of salt as the deciding factor.

    Like

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