Notices
Results 1 to 26 of 26

Thread: Sodium Carbonate

  1. #1 Sodium Carbonate 
    New Member
    Join Date
    Jan 2018
    Posts
    1
    I put approximately 6 cups of a sodium Carbonate based laundry powder into a large aluminum canning pot with about 10 gallons of water. I placed some metal (mostly steel I believe) objects in it and boiled it for an hour. Intent was to degrease in prep for painting. Now I have a fairly heavy coat of chalky film. It's hard stuff to scrub and scrape as there are springs and chains. What procedure would "defilm" these items? Also I have some calcified deposits in the pot itself. Is this corrosion? What can be done? Thank you in advance for your time!


    Reply With Quote  
     

  2.  
     

  3. #2  
    Bullshit Intolerant PhDemon's Avatar
    Join Date
    Feb 2013
    Location
    Newcastle-upon-Tyne, UK
    Posts
    4,909
    Sodium carbonate will break down to sodium oxide on heating. This should be water soluble but if it is clogged on may need scrubbing. Try soaking the stuff in vinegar, this will get rid of the sodium oxide and also any insoluble calcium salts from hard water.


    Reply With Quote  
     

  4. #3  
    Bullshit Intolerant PhDemon's Avatar
    Join Date
    Feb 2013
    Location
    Newcastle-upon-Tyne, UK
    Posts
    4,909
    Also for degreasing an organic solvent may have been better than sodium carbonate solution... Dichloromethane if you can get it would be worth a try, if not an alkane like cyclohexane may work.

    (Benzene would be the solvent of choice but due to its carcinogenic properties it's no longer that easy to get hold of).
    Reply With Quote  
     

  5. #4  
    exchemist
    Join Date
    May 2013
    Location
    London
    Posts
    2,921
    Quote Originally Posted by PhDemon View Post
    Sodium carbonate will break down to sodium oxide on heating. This should be water soluble but if it is clogged on may need scrubbing. Try soaking the stuff in vinegar, this will get rid of the sodium oxide and also any insoluble calcium salts from hard water.
    Could there not be more going on here? If the pot is Al, then won't Na2CO3 be alkaline enough to attack the passivating oxide layer, creating aluminate ions?
    Or is my memory faulty?
    Reply With Quote  
     

  6. #5  
    Bullshit Intolerant PhDemon's Avatar
    Join Date
    Feb 2013
    Location
    Newcastle-upon-Tyne, UK
    Posts
    4,909
    Possibly, the reaction of aluminium (III) ions with carbonate will produce aluminium hydroxide which is a precipitate. As it is an insoluble hydroxide I think reacting it with vinegar would still remove it. You would need to add an excess of sodium hydroxide to form the aluminate(III) I don't think the carbonate can do this, and in any case the aluminate is soluble in water.
    Reply With Quote  
     

  7. #6  
    exchemist
    Join Date
    May 2013
    Location
    London
    Posts
    2,921
    Quote Originally Posted by PhDemon View Post
    Possibly, the reaction of aluminium (III) ions with carbonate will produce aluminium hydroxide which is a precipitate. As it is an insoluble hydroxide I think reacting it with vinegar would still remove it. You would need to add an excess of sodium hydroxide to form the aluminate(III) I don't think the carbonate can do this, and in any case the aluminate is soluble in water.
    Ah OK.
    Reply With Quote  
     

  8. #7  
    Forum Isotope
    Join Date
    Feb 2012
    Location
    Western US
    Posts
    2,794
    I have a question for you experts. In my youth I enjoyed making nitrogen iodide "poppers" (although the iodine stains were not easily cleaned). I recall one of my chemistry teachers commenting that this substance was the only chemical that could be triggered into detonation by alpha rays. Is this in fact true? How does nitrogen trichloride compare in this regard (I know it's sufficiently more dangerous that my teacher looked horrified when I expressed a desire to make some)?
    Reply With Quote  
     

  9. #8  
    Bullshit Intolerant PhDemon's Avatar
    Join Date
    Feb 2013
    Location
    Newcastle-upon-Tyne, UK
    Posts
    4,909
    I'm not sure about the alpha ray fact but it could be true as the N-I bond is quite weak (~169 kJ/mol), most other explosives are a bit more stable with stronger covalent bonds.

    As to how NCl3 compares to NI3 the amount of bang you get for your buck is determined by the relative strengths of the N-X and X-X bonds. I've done a back of the envelope calculation and the enthalpy change when NCl3 detonates is -232 kJ/mol, when NI3 detonates it is ~-150 kJ/mol. So you are getting 1.5 times more energy output per mole of NCl3 as compared to NI3. If you factor in the molar mass, NCl3 produces 5 times more energy than NI3 gram for gram.

    Also NCl3 is a liquid whereas NI3 is a solid, solid explosives tend to be easier to handle. Also NCl3 is more toxic than NI3 and produces Cl2 when it goes off, this is nastier than the iodine you get from the iodide.

    All of these may have contributed to your teacher's horror!

    Thanks for making me think about this. I'm going to set this as a problem when my A level chemistry class revise energetics and thermodynamics
    Reply With Quote  
     

  10. #9  
    exchemist
    Join Date
    May 2013
    Location
    London
    Posts
    2,921
    Quote Originally Posted by PhDemon View Post
    I'm not sure about the alpha ray fact but it could be true as the N-I bond is quite weak (~169 kJ/mol), most other explosives are a bit more stable with stronger covalent bonds.

    As to how NCl3 compares to NI3 the amount of bang you get for your buck is determined by the relative strengths of the N-X and X-X bonds. I've done a back of the envelope calculation and the enthalpy change when NCl3 detonates is -232 kJ/mol, when NI3 detonates it is ~-150 kJ/mol. So you are getting 1.5 times more energy output per mole of NCl3 as compared to NI3. If you factor in the molar mass, NCl3 produces 5 times more energy than NI3 gram for gram.

    Also NCl3 is a liquid whereas NI3 is a solid, solid explosives tend to be easier to handle. Also NCl3 is more toxic than NI3 and produces Cl2 when it goes off, this is nastier than the iodine you get from the iodide.

    All of these may have contributed to your teacher's horror!

    Thanks for making me think about this. I'm going to set this as a problem when my A level chemistry class revise energetics and thermodynamics
    The Wiki article says that the notorious instability of NI3 is due to the steric strain imposed by the large size of the 3 I atoms, leading to a very low activation energy for the decomposition. This would suggest that NCl3 might be a lot more kinetically stable, even if thermodynamically you get more energy released from it. (Trick question potential here?)

    I was not aware of the alpha particle thing either, thought Wiki mentions it. I suppose capture or rebound of an alpha particle by N or I would provide enough momentum to break one bond in one molecule and, if the activation barrier is low enough, this could make initiation of a chain reaction more probable than other dissipative processes for the radicals (or ions?) created.

    Also interesting that the structure when made in the lab is an adduct NI3.NH3 (this much I had remembered!), consisting apparently of chains of NI2-I-NI2-I-.......with NH3 molecules between the chains.

    It's actually a great discussion topic in lots of ways - and the little puffs of purple smoke are so pretty!
    Reply With Quote  
     

  11. #10  
    Forum Isotope
    Join Date
    Feb 2012
    Location
    Western US
    Posts
    2,794
    Quote Originally Posted by exchemist View Post
    Quote Originally Posted by PhDemon View Post
    I'm not sure about the alpha ray fact but it could be true as the N-I bond is quite weak (~169 kJ/mol), most other explosives are a bit more stable with stronger covalent bonds.

    As to how NCl3 compares to NI3 the amount of bang you get for your buck is determined by the relative strengths of the N-X and X-X bonds. I've done a back of the envelope calculation and the enthalpy change when NCl3 detonates is -232 kJ/mol, when NI3 detonates it is ~-150 kJ/mol. So you are getting 1.5 times more energy output per mole of NCl3 as compared to NI3. If you factor in the molar mass, NCl3 produces 5 times more energy than NI3 gram for gram.

    Also NCl3 is a liquid whereas NI3 is a solid, solid explosives tend to be easier to handle. Also NCl3 is more toxic than NI3 and produces Cl2 when it goes off, this is nastier than the iodine you get from the iodide.

    All of these may have contributed to your teacher's horror!

    Thanks for making me think about this. I'm going to set this as a problem when my A level chemistry class revise energetics and thermodynamics
    The Wiki article says that the notorious instability of NI3 is due to the steric strain imposed by the large size of the 3 I atoms, leading to a very low activation energy for the decomposition. This would suggest that NCl3 might be a lot more kinetically stable, even if thermodynamically you get more energy released from it. (Trick question potential here?)

    I was not aware of the alpha particle thing either, thought Wiki mentions it. I suppose capture or rebound of an alpha particle by N or I would provide enough momentum to break one bond in one molecule and, if the activation barrier is low enough, this could make initiation of a chain reaction more probable than other dissipative processes for the radicals (or ions?) created.

    Also interesting that the structure when made in the lab is an adduct NI3.NH3 (this much I had remembered!), consisting apparently of chains of NI2-I-NI2-I-.......with NH3 molecules between the chains.

    It's actually a great discussion topic in lots of ways - and the little puffs of purple smoke are so pretty!
    Thank you, exchemist and PhDemon -- I knew that I was going to learn a great deal from your answers, and you did not disappoint!

    So, the iodide triggers more easily, but one gets a much bigger bang from the trichloride. The iodide leaves a pretty stain, while the trichloride singes your lungs. Excellent.

    A little googling turns up many tales of danger from the trichloride. Apparently Dulong was the first to publish, at the cost of a couple of fingers and an eye. He omitted mention of his injuries, so that Davy nearly lost his sight when trying to replicate the experiment. I now understand the look of horror on my teacher's face!
    Reply With Quote  
     

  12. #11  
    exchemist
    Join Date
    May 2013
    Location
    London
    Posts
    2,921
    Quote Originally Posted by tk421 View Post
    Quote Originally Posted by exchemist View Post
    Quote Originally Posted by PhDemon View Post
    I'm not sure about the alpha ray fact but it could be true as the N-I bond is quite weak (~169 kJ/mol), most other explosives are a bit more stable with stronger covalent bonds.

    As to how NCl3 compares to NI3 the amount of bang you get for your buck is determined by the relative strengths of the N-X and X-X bonds. I've done a back of the envelope calculation and the enthalpy change when NCl3 detonates is -232 kJ/mol, when NI3 detonates it is ~-150 kJ/mol. So you are getting 1.5 times more energy output per mole of NCl3 as compared to NI3. If you factor in the molar mass, NCl3 produces 5 times more energy than NI3 gram for gram.

    Also NCl3 is a liquid whereas NI3 is a solid, solid explosives tend to be easier to handle. Also NCl3 is more toxic than NI3 and produces Cl2 when it goes off, this is nastier than the iodine you get from the iodide.

    All of these may have contributed to your teacher's horror!

    Thanks for making me think about this. I'm going to set this as a problem when my A level chemistry class revise energetics and thermodynamics
    The Wiki article says that the notorious instability of NI3 is due to the steric strain imposed by the large size of the 3 I atoms, leading to a very low activation energy for the decomposition. This would suggest that NCl3 might be a lot more kinetically stable, even if thermodynamically you get more energy released from it. (Trick question potential here?)

    I was not aware of the alpha particle thing either, thought Wiki mentions it. I suppose capture or rebound of an alpha particle by N or I would provide enough momentum to break one bond in one molecule and, if the activation barrier is low enough, this could make initiation of a chain reaction more probable than other dissipative processes for the radicals (or ions?) created.

    Also interesting that the structure when made in the lab is an adduct NI3.NH3 (this much I had remembered!), consisting apparently of chains of NI2-I-NI2-I-.......with NH3 molecules between the chains.

    It's actually a great discussion topic in lots of ways - and the little puffs of purple smoke are so pretty!
    Thank you, exchemist and PhDemon -- I knew that I was going to learn a great deal from your answers, and you did not disappoint!

    So, the iodide triggers more easily, but one gets a much bigger bang from the trichloride. The iodide leaves a pretty stain, while the trichloride singes your lungs. Excellent.

    A little googling turns up many tales of danger from the trichloride. Apparently Dulong was the first to publish, at the cost of a couple of fingers and an eye. He omitted mention of his injuries, so that Davy nearly lost his sight when trying to replicate the experiment. I now understand the look of horror on my teacher's face!
    Something I did not know. Thanks.

    In fact, I was prompted by this exchange to look up "phlegmatisers" - substances that are added to sensitive explosives to make them safer to handle. I suppose the most famous of these was kieselguhr (diatomaceous earth) used for nitroglycerine, in dynamite. What I am having trouble finding is a mechanistic explanation of the mode of action of these phlegmatisers. I presume they offer a wider variety of deactivation pathways to those molecules that acquire enough energy to disintegrate, but would quite like to see more about the actual competing processes.
    Last edited by exchemist; February 22nd, 2018 at 12:57 PM.
    Reply With Quote  
     

  13. #12  
    Bullshit Intolerant PhDemon's Avatar
    Join Date
    Feb 2013
    Location
    Newcastle-upon-Tyne, UK
    Posts
    4,909
    I think Faraday was also injured by it, when Davy was indisposed due to NCl3 detonation he got Faraday to carry on the experiments! What a lab supervisor! Pre-COSHH and health and safety obviously!
    Reply With Quote  
     

  14. #13  
    Forum Isotope
    Join Date
    Feb 2012
    Location
    Western US
    Posts
    2,794
    Quote Originally Posted by PhDemon View Post
    I think Faraday was also injured by it, when Davy was indisposed due to NCl3 detonation he got Faraday to carry on the experiments! What a lab supervisor! Pre-COSHH and health and safety obviously!
    So, among his many firsts, Faraday was the first graduate student to be hired as cannon fodder? Brilliant!
    Reply With Quote  
     

  15. #14  
    Forum Isotope
    Join Date
    Feb 2012
    Location
    Western US
    Posts
    2,794
    Quote Originally Posted by exchemist View Post
    In fact, I was prompted by this exchange to look up "phegmatisers" - substances that are added to sensitive explosives to make them safer to handle. I suppose the most famous of these was kieselguhr (diatomaceous earth) used for nitroglycerine, in dynamite. What I am having trouble finding is a mechanistic explanation of the mode of action of these phlegmatisers. I presume they offer a wider variety of deactivation pathways to those molecules that acquire enough energy to disintegrate, but would quite like to see more about the actual competing processes.
    Thanks for teaching me about "phlegmatisers" -- I'd not come across that term before. I'd always wondered about the mechanism behind kieselguhr's moderating action. In histories of Nobel I can't recall ever coming across an explanation of how kieselguhr makes dynamite more phlegmatic. I wonder if he had a specific mechanism in mind, or if he took an Edisonian approach to the problem. It appears I have some reading to do. Thank you, exchemist!
    Reply With Quote  
     

  16. #15  
    Bullshit Intolerant PhDemon's Avatar
    Join Date
    Feb 2013
    Location
    Newcastle-upon-Tyne, UK
    Posts
    4,909
    Also, the iodine stains can easily be cleaned with sodium thiosulphate solution (old school photographers "hypo"). This will reduce the iodine to iodide which is water soluble and can be wiped up with a damp cloth.
    Reply With Quote  
     

  17. #16  
    Forum Isotope
    Join Date
    Feb 2012
    Location
    Western US
    Posts
    2,794
    Quote Originally Posted by PhDemon View Post
    Also, the iodine stains can easily be cleaned with sodium thiosulphate solution (old school photographers "hypo"). This will reduce the iodine to iodide which is water soluble and can be wiped up with a damp cloth.
    That's probably what my chemistry teacher gave me to clean up the mess I made. Thanks for letting me know. Maybe it's time to make some more "poppers" and relive an episode from my youth...
    Reply With Quote  
     

  18. #17  
    exchemist
    Join Date
    May 2013
    Location
    London
    Posts
    2,921
    Quote Originally Posted by tk421 View Post
    Quote Originally Posted by exchemist View Post
    In fact, I was prompted by this exchange to look up "phegmatisers" - substances that are added to sensitive explosives to make them safer to handle. I suppose the most famous of these was kieselguhr (diatomaceous earth) used for nitroglycerine, in dynamite. What I am having trouble finding is a mechanistic explanation of the mode of action of these phlegmatisers. I presume they offer a wider variety of deactivation pathways to those molecules that acquire enough energy to disintegrate, but would quite like to see more about the actual competing processes.
    Thanks for teaching me about "phlegmatisers" -- I'd not come across that term before. I'd always wondered about the mechanism behind kieselguhr's moderating action. In histories of Nobel I can't recall ever coming across an explanation of how kieselguhr makes dynamite more phlegmatic. I wonder if he had a specific mechanism in mind, or if he took an Edisonian approach to the problem. It appears I have some reading to do. Thank you, exchemist!
    Still no luck with the phlegmatiser mechanism, I'm afraid.

    However I have had a further thought regarding the energy release. PhDemon did an enthalpy calculation suggesting decomposition of NCl3 would be more exothermic than NI3. I wonder if his calculation would have taken account of the steric strain imposed by 3 large I atoms around N. I imagine NI3 may be forced into a nearly planar (sp2) configuration and that the effective bond energies may be greatly reduced by this, leading to a more exothermic reaction that would be predicted from using 3 times the standard single N-I bond energy.

    In fact I see that, looking at the Wiki article, it says the decomposition yields 290kJ/mol rather than the 150kJ/mol that he calculated, which seems to bear out the idea that steric strain has a seriously destabilising effect on the molecule. However I can't find a bond angle for NI3, which would be nice to know. Maybe nobody dares put this stuff into an expensive X-Ray crystallography apparatus!

    Anyway, possibly yet another interesting point for 6th form discussion: how steric effects complicate bond energy calculations!

    What an interesting molecule. Surely there must be, somewhere, a PhD chemist (with one eye and a couple of missing fingers) who has researched all this.....
    Reply With Quote  
     

  19. #18  
    Bullshit Intolerant PhDemon's Avatar
    Join Date
    Feb 2013
    Location
    Newcastle-upon-Tyne, UK
    Posts
    4,909
    My calculation just included bond enthalpies, I took no account of steric factors...

    The only bond angle information I could find was here

    https://www.google.co.uk/amp/s/amp.r..._why_nitrogen/

    Which says 101-102 degrees but that is for the ammonia adduct... I'll check if it is in my CRC Handbook tomorrow and do a more detailed search tomorrow.
    Reply With Quote  
     

  20. #19  
    exchemist
    Join Date
    May 2013
    Location
    London
    Posts
    2,921
    Quote Originally Posted by PhDemon View Post
    My calculation just included bond enthalpies, I took no account of steric factors...

    The only bond angle information I could find was here

    https://www.google.co.uk/amp/s/amp.r..._why_nitrogen/

    Which says 101-102 degrees but that is for the ammonia adduct... I'll check if it is in my CRC Handbook tomorrow and do a more detailed search tomorrow.
    Blimey! Good point - I was forgetting about the chain structure -I-NI2-I-NI2- . So N has a coordination number of 4 in this structure, with 2 different environments for I atoms. Even more complicated. Presume there will thus be several different applicable bond angles.
    Reply With Quote  
     

  21. #20  
    Bullshit Intolerant PhDemon's Avatar
    Join Date
    Feb 2013
    Location
    Newcastle-upon-Tyne, UK
    Posts
    4,909
    I couldn't find any further info, but MolCalc (see link below) calculates a bond angle of 101.4 degrees...


    Molecule Calculator (MolCalc)

    to calculate the bond angle double click on an I atom, single click on the N and double click on a different I atom)
    Reply With Quote  
     

  22. #21  
    exchemist
    Join Date
    May 2013
    Location
    London
    Posts
    2,921
    Quote Originally Posted by PhDemon View Post
    I couldn't find any further info, but MolCalc (see link below) calculates a bond angle of 101.4 degrees...


    Molecule Calculator (MolCalc)

    to calculate the bond angle double click on an I atom, single click on the N and double click on a different I atom)
    And that is for NI3! I find that very counterintuitive. I was thinking steric strain would flatten the molecule, opening the angle further than the regular tetrahedral angle (109.5?), but it seems to calculate the opposite. Why would that be?
    Reply With Quote  
     

  23. #22  
    exchemist
    Join Date
    May 2013
    Location
    London
    Posts
    2,921
    Quote Originally Posted by PhDemon View Post
    I couldn't find any further info, but MolCalc (see link below) calculates a bond angle of 101.4 degrees...


    Molecule Calculator (MolCalc)

    to calculate the bond angle double click on an I atom, single click on the N and double click on a different I atom)
    And that is for NI3! I find that very counterintuitive. I was thinking steric strain would flatten the molecule, opening the angle further than the regular tetrahedral angle (109.5?), but it seems to calculate the opposite. Why would that be?
    Reply With Quote  
     

  24. #23  
    Bullshit Intolerant PhDemon's Avatar
    Join Date
    Feb 2013
    Location
    Newcastle-upon-Tyne, UK
    Posts
    4,909
    Top of the head after a few beers but...

    The N-I bond is not very polar, there will be more electron density around the lone pair on nitrogen than in the chloride. Lone pair-bonding pair repulsion is greater than bonding pair-bonding pair, so unless the I atoms actually touch the lone pair repulsion wins reducing the bond angle.
    Reply With Quote  
     

  25. #24  
    Forum Isotope
    Join Date
    Feb 2012
    Location
    Western US
    Posts
    2,794
    Quote Originally Posted by PhDemon View Post
    Top of the head after a few beers but...

    The N-I bond is not very polar, there will be more electron density around the lone pair on nitrogen than in the chloride. Lone pair-bonding pair repulsion is greater than bonding pair-bonding pair, so unless the I atoms actually touch the lone pair repulsion wins reducing the bond angle.
    Extremely interesting. You guys are giving me a terrific education. Thank you very much!
    Reply With Quote  
     

  26. #25  
    Bullshit Intolerant PhDemon's Avatar
    Join Date
    Feb 2013
    Location
    Newcastle-upon-Tyne, UK
    Posts
    4,909
    Thanks for the kind words but I got it arse about face (the problems of posting from the pub!)

    The nitrogen atom in NI3 does have more electron density than in NCl3 so the argument stands but my statement "the N-I bond is not very polar", was ex ano!
    Reply With Quote  
     

  27. #26  
    Forum Isotope
    Join Date
    Feb 2012
    Location
    Western US
    Posts
    2,794
    Quote Originally Posted by PhDemon View Post
    Thanks for the kind words but I got it arse about face (the problems of posting from the pub!)

    The nitrogen atom in NI3 does have more electron density than in NCl3 so the argument stands but my statement "the N-I bond is not very polar", was ex ano!
    Not a problem, as I had had a pint myself, so it all made sense.
    Reply With Quote  
     

Similar Threads

  1. preparation of sodium carbonate
    By diala9721 in forum Chemistry
    Replies: 2
    Last Post: February 10th, 2014, 06:56 AM
  2. Replies: 3
    Last Post: January 6th, 2014, 06:30 AM
  3. Hydrogen Carbonate reactions....
    By 09jnewington in forum Chemistry
    Replies: 1
    Last Post: November 25th, 2013, 01:50 PM
  4. Get sodium hydroxide from sodium-metabisulphite
    By serbianchemist in forum Chemistry
    Replies: 9
    Last Post: June 11th, 2012, 05:39 PM
  5. sodium hydrogen carbonate concentration,
    By A.londoner in forum Biology
    Replies: 2
    Last Post: September 4th, 2006, 08:26 AM
Bookmarks
Bookmarks
Posting Permissions
  • You may not post new threads
  • You may not post replies
  • You may not post attachments
  • You may not edit your posts
  •