Sciencemadness Discussion Board

HCL synthesis

cranium - 8-8-2006 at 18:29

I produced what I believe is HCL in solution by mixing 2mL hydrogen peroxide, 2mL citric acid, and 10mg sodium chloride. The product of the reaction has a pH of 1.5. I believe it is HCl because the H+ ions in the acid most likely bonded with the Cl- ions in the sodium chloride. What do you guys think was produced? It works great as a cleaning solution to remove glue that is caked on and almost impossible to remove after you work with polymers, yet is safe to touch (for the record, I am aware of the dangers of HCL, lol):P.

guy - 8-8-2006 at 18:50

There's no way it could have produced HCl. HCl is a stronger acid than citric acid so it would make citrate ions in to citric acid again.

BromicAcid - 8-8-2006 at 18:59

If the solution is acidic then there are hydronium ions, if there are also chloride ions present then technically you can say there is HCl present in the solution. Really though the acidity is just due to the citric acid.

Acidic peroxide soluitons are good cleaning agents, acetic acid and H<sub>2</sub>O<sub>2</sub> is a well known cleaning agent. The use of salt is also sometimes employed, usually in excess to act as an abrasive.

What is the concentration of peroxide and citric acid used? Citric acid can act as a reducing agent and hydrogen peroxide an oxidizing agent so there can be interactions there, which may be expecailly quick if the concentrations are high enough.

fescho - 12-8-2006 at 09:17

think the same... ion bond between NaCl is too strong to be broken just by citric acid or hydrogen peroxide... and Na+ is strong base which is there neutralised by Cl- so the pH of 1.5 is not made by Cl...

[Edited on 12-8-2006 by fescho]

Nicodem - 12-8-2006 at 11:37

The bond in NaCl strong? You must be joking, I break that bond on daily basis.:D All it takes is some water. You dissolve NaCl in water and you have Na+ and Cl- safely surrounded by H2O molecules. Compared to covalent bonds, ionic interactions are quite weak bonds (if they can be called bonds at all!).

Na+ is not a base and Cl- is not an acid, actually the chloride anion is a very weak base. Read on acids and bases first.

guy - 12-8-2006 at 12:21

Quote:
Originally posted by Nicodem
The bond in NaCl strong? You must be joking, I break that bond on daily basis.:D All it takes is some water. You dissolve NaCl in water and you have Na+ and Cl- safely surrounded by H2O molecules. Compared to covalent bonds, ionic interactions are quite weak bonds (if they can be called bonds at all!).

Na+ is not a base and Cl- is not an acid, actually the chloride anion is a very weak base. Read on acids and bases first.


Actually ionic "bonds" are strong but it is overcome due to hydration energy and entropy.

fescho - 12-8-2006 at 23:44

yeah but can you break it without dissolving in water? can by electrolysis at temper around 700*C... an i didn´t said that Cl is an acid, but Cl- ions of HCl will be react with Na+... don´t understand how can exist a mixture of natrium citric and HCl... of course if it would exist... and can agree with you, i´m not expert in chem but i wanna be...:P

and thanks for that link... i´ll check it...

and when you solve NaCl in water you don´t break a ˇbondˇ... you just put Na+ and Cl- from each other... the bond still exist cause hydratated are ions not atoms like Na and Cl... don´t know it´s right but think :)

[Edited on 13-8-2006 by fescho]

Nicodem - 13-8-2006 at 01:35

Quote:
Originally posted by fescho
yeah but can you break it without dissolving in water?

But that is exactly what you don’t understand well. There is no bond in purely ionic crystals and NaCl is such an ionic compound. Since there is no bond between Na+ and Cl+ it is also hard to talk about what is its heterolytic bond dissociation energy. As Guy already mentioned, it is easy to counteract the coulumbic (electrostatic) forces that holds an ionic crystal together by solvatation in dielectric solvents (like the mentioned hydratation in H2O). The energy needed to separate ions is provided by solvatation and the coulumbic interaction is inhibited by the dielectricity of the solvent. But the energy needed for heterolytic bond dissociation energy of a true bond, like a covalent bond, is quite higher and indeed a covalent bond will much prefer to dissociate in a homolytic manner where a smaller energy is needed to break a bond (since neutral radicals form there is no coulumbic force to counteract). In solution chemistry there are exceptions like acids which rather dissociate heterolyticaly in dielectric solvents (like HCl in H2O).
Quote:
Originally posted by fescho
and when you solve NaCl in water you don´t break a ˇbondˇ... you just put Na+ and Cl- from each other... the bond still exist cause hydratated are ions not atoms like Na and Cl... don´t know it´s right but think


Your concepts are confused. NaCl is not a compound of Na and Cl but a crystal composed of Na+ and Cl- ions. These two ions are already apart from each other by the distance of the crystalline lattice. Dissolving the crystal in water only breaks these crystal lattices and its organized structure and thus the ions are free to move around in solution. So here is no real bond to break in the first place, there is only the strong Coulomb interaction that has to be counteracted.

For example, when you dissolve sodium chloride, citric acid and hydrogen peroxide in water you will have a water solution of several ions and several species in equilibrium:

NaCl(s) <=> Na(+)(aq) + Cl(-)(aq)
(determined by the solubility equilibrium constant)

H3[citrate] <=> H(+) + H2[citrate](-)
H2[citrate](-) <=> H(+) + H[citrate](2-)
H[citrate](2-) <=> H(+) + [citrate](3-)
(the three dissociation steps of citric acid, determined by pKa1=3.15, pKa2=4.77 and pKa3=6.40 acidity constants)

H2O2 <=> H(+) + HOO(-) (hydrogen peroxide is a weak acid, pKa1=11.65, so this equilibrium is almost negligible in acidic media)

Same goes for the following reactions with very low equilibrium constant:
H2O <=> H(+) + OH(-) (the classical self dissociation of water)
H(+) + Cl(-) <=> HCl (so actually HCl does form as you suspected, but in negligible amounts as the pKa of HCl is -8.0 so the equilibrium is way to the left even at pH as low as 1.5)

These are only ionic reactions and even not all of the possible ones. If you consider also the redox reactions you can see there are several possibilities of H2O2 getting involved in the oxidation of citric acid, citrate anions, chloride anions and who knows what else.

PS: Read the provided links thoroughly. I hate search for references for beginners when they rather stubbornly stuck with they wrong concepts.

woelen - 13-8-2006 at 04:07

Fescho, I have to say that nicodem is on the right track. NaCl is not a simple compound like H2O or citric acid.

Nicodem has explained things quite well, so I will not repeat that. I only want to add another point to make things clear.

Suppose you have 1 mol of NaBr and 1 mol of KCl and you add this to one liter of water. All of both solids dissolves. Label this liquid A.

Now you also have 1 mol of NaCl and 1 mol of KBr and you add this to one liter of water. Again, all of both solids dissolved. Label this liquid B.

Now you give solutions A and B to an expert chemist and ask him to tell you which is the NaBr+KCl and which is the NaCl+KBr. Even the best chemist of the world, supplied with the best apparatus in the world will not be able to tell you which is solution A and which is solution B. They are completely the same. This is because in solution you cannot speak of NaCl, NaBr, KCl and KBr, but you have Na(+), K(+), Cl(-) and Br(-) ions in solution, all in a 1:1:1:1 ratio.

fescho - 13-8-2006 at 06:34

Quote:
H(+) + Cl(-) <=> HCl (so actually HCl does form as you suspected, but in negligible amounts as the pKa of HCl is -8.0 so the equilibrium is way to the left even at pH as low as 1.5)


i´m really sorry but this i don´t understand... thanks for lot of links... i´ll check them soon and thanks for explaining and sorry for bother you with this...

and i have still one question... is this reaction real this?

HOO(-) <=> H(+) + OO(2-)
cuz HOO(-) is still acid and i don't know if it exist...

thanks...;)

[Edited on 13-8-2006 by fescho]

Nicodem - 13-8-2006 at 08:30

The acid dissociation constant (pKa) of HCl is very low (HCl is a strong acid) which means that when you dissolve HCl in water there will be very little undissociated HCl left.
For example, let’s say you dissolve 0.1 mol of HCl in 1L of water. You might call this a 0.1 mol/l aqueous solution of HCl, but in reality you will have a solution of approximately 0.1 mol/l chloride anions (Cl-), 0.1 mol/l hydronium cations (H3O+) and only a negligible concentration of HCl.
You can estimate the concentration of HCl in such a solution by the use of the pKa value of HCl, which is the negative decimal logarithm of the reaction equilibrium constant (the operator p is used for convenience only, since it is easier to write pKa = -8 rather than Ka = 100000000).

The relevant reaction is that of water protonation by HCl (note that H2O acts as a base here!):

HCl + H2O <=> H3O(+) + Cl(-)

The reaction equilibrium constant is approximately:

Ka = 10^(-pKa) = 100000000

(it is hard to exactly measure pKa of strong acids thus various sources give different values for HCl, here we took it to be pKa = -8 as reported in Wikipedia)

The equilibrium constant equation is:

Ka = ( [H3O(+)] × [Cl(-)] ) / [HCl]

where [X] represents the activity of the species X. Since we look only for an estimation we take activity to equal concentration. Also, since HCl is a strong acid and our solution is diluted we are allowed to simplify the calculation by assuming near total dissociation (the error will be less than the error derived from the inaccuracy of pKa value). Thus:

[H3O(+)] = 0.1 mol/l (which means the pH of the solution will approximately be pH = –log(0.1) = 1)
and
[Cl(-)] = 0.1 mol/l

Now we have all we need to calculate [HCl]:

[HCl] = ( [H3O(+)] × [Cl(-)] ) / Ka = ( 0.1 mol/l × 0.1 mol/l ) / ( 0.0000001 mol/l ) = 0.0000000001 mol /l

So as you see the concentration of HCl in this solution is nearly negligible, only 10^(-10) mol/l!

For your mixture of NaCl, H2O2 and citric acid with the measured pH of 1.5 you can calculate yourself the estimated concentration of HCl and report back the results. You will need the concentration of NaCl though. The [H3O(+)] can be calculated from the pH value. When you will obtain the concentration of the formed HCl you will understand why already BromicAcid told you that “technically you can say there is HCl present in the solution”.
Quote:

and i have still one question... is this reaction real this?

HOO(-) <=> H(+) + OO(2-)
cuz HOO(-) is still acid and i don't know if it exist...


The pKa2 of H2O2 is the value that could answer that. As far as I know, it was never measured and is probably pretty hard to measure (I never saw it anywhere). The O2(2-) anion is a very strong base and the equilibrium constant of its formation in aqueous solutions could only be estimated by comparison with other strong bases. For example, you can safely deduce that O2(2-) it is somewhat less basic than O(2-), but in any case I would say the value of pKa of its conjugated acid is more than 30. Thus you can safely assume that the probability of such dissociation in water is near to zero. As such, the significance of that equilibrium that you wrote is also zero in any kind of conceivable reactions that you might assume involving O2(2-). For example, lets assume you have a reaction where H2O2 precipitates a divalent metal cation M(2+) in the form of MO2. You might argue that the precipitation reaction is:

M(2+)(aq) + O2(2-)(aq) <=> MO2(s)

and formaly that could be correct but the true mechanism of the precipitation need not to call for the existance of O2(2-) in the solution. It would more likely be like this:

M(2+)(aq) + HOO(-)(aq) <=> HOOM(+)(aq) <=> MO2(s) + H(+)(aq)

Do you understand this? I would say it is a bit too much for secondary school, so perhaps you should read books instead.

fescho - 15-8-2006 at 10:06

thanks for explaining... i understand now all except that pKa... i have never heard it... so few more question :)

pKa = -log Ka

Ka = 10e-pKa ?

and thanks for time you spent to answer...
i know it all just in Slovak... cuz i haven´t found books with basics of chem in English... please can you recommend me some? i´m missing the terminology and lots of other... really thanks...

YT2095 - 15-8-2006 at 12:20

Quote:
Originally posted by cranium
I produced what I believe is HCL in solution by mixing 2mL hydrogen peroxide, 2mL citric acid, and 10mg sodium chloride.


How did you get 2ml of Citric acid?:P

ethan_c - 15-8-2006 at 13:55

Quote:
Originally posted by YT2095
Quote:
Originally posted by cranium
I produced what I believe is HCL in solution by mixing 2mL hydrogen peroxide, 2mL citric acid, and 10mg sodium chloride.


How did you get 2ml of Citric acid?:P


He squeezed 2mL out of a lemon, silly. Everyone knows that's how you get citric acid, straight from citrus fruit!

Nicodem - 16-8-2006 at 10:48

Cranium was obviously just trolling like in all the others threads he started under this as well as his previus nickname. However, I wander why many like to reply and spur him so much? Well, I guess it's some kind of fun after all.


To Fescho:
The pKa value is a specific property of any compound that can dissociate to protons. Thus it has to be measured.
It tells how strong the acid is or put in other words it tells where the equilibrium of the dissociation reaction lies (like you already understood, it is the negative decimal logarithm of the equilibrium constant). It is also highly specific to the solvent used in the measurement as well as the temperature (just as any reaction equilibrium).
The lower the value the strongest the acid (HClO4 is one of the strongest acids and has the pKa of about -10). Weak acids have high values of pKa. For example water is a weak acid and its pKa is about 15. Ammonia (NH3) is one of the weakest acids having the pKa of 35. On the other side the anions of weak acids make for strong bases. For example, the anion you get from the dissociation of H2O is the hydroxide (OH-). Thus solutions of hydroxides are basic. The anion you get from ammonia dissociation is the amide (H2N-) and sodium amide (NaNH2) is one of the strongest bases. The pKa value can thus be used for bases as well when it is understood that one is talking about the conjugate acid of the base. But in order to avoid confusion there is also the pKb value for bases. One can calculate the pKb from the pKa value of the conjugate acid if the self dissociation constant of the solvent is known (for water as a solvent: pKa+pKb=pKw).

Also read:
http://en.wikipedia.org/wiki/Acids_and_bases
http://en.wikipedia.org/wiki/Acid_dissociation_constant

Tables with pKa values of many common acids in water solution:
http://research.chem.psu.edu/brpgroup/pKa_compilation.pdf
http://www.its.caltech.edu/~travisw/pkas.pdf (a smaller one with organics mainly)