Sciencemadness Discussion Board

Transition metal nitrites

symboom - 15-7-2018 at 01:29



Titanium iodide and nitrogen dioxide forms
Titanium nitrite and iodine

Vandium iodide and nitrogen dioxide forms
Vandium nitrite and iodine

Chromium sulfate and sodium nitrite
Chromium nitrite
................................
Properties of sodium nitrite
soluble in methanol (4.4 g/100 mL)
ethanol
slightly soluble in diethyl ether (0.3 g/100 mL)
very soluble in ammonia


This is my results

Cobalt nitrite attempt
Cobalt nitrate and sodium nitrite
Deep red souble compound
Red compound distroyed by NaClO

manganese nitrite attempt
Manganese sulfate and sodium nitrite
Pink souble compound destroyed with NaClO
No addition of H2O2 catalytic effect

Nivkel sulfate is added to a solution of sodium nitrite
Nickel nitrite forms deep green color is observed then decomposes in H2O


A solution of Nickel Nitrite, Ni(NO2)2, is obtained by double decomposition of nickel sulphate and barium nitrite, but the salt itself has not been isolated. It yields a stable double salt with potassium nitrite, namely, 4KNO2.Ni(NO2)2, which crystallises in brownish red octahedra, when excess of potassium nitrite is added to a concentrated solution of nickel nitrite. In the presence of a calcium salt a yellow crystalline precipitate is obtained, of composition represented by the formula 2KNO2.Ca(NO2)2.2Ni(NO2)2 or K2CaNi(NO2)6, which closely resembles potassium cobalti-nitrite in appearance. The corresponding barium and strontium salts have been prepared.

Copper nitrite
Copper sulfate is added to a solution of sodium nitrite
Copper nitrite forms a deep green color is observed then decomposes in H2O
My Hypothysis is to Prepare in alcohol

The nitrite is only known in solution, prepared by addition of lead nitrite to cupric-sulphate solution. Exposure of its dilute solution to air causes slow formation of nitrate. On evaporation of a concentrated solution over sulphuric acid, there is partial decomposition in accordance with the equation

3Cu(NO2)2 = Cu(NO3)2 + 2CuO + 4NO.

It forms a number of complex nitrites with other metals.


Zinc
Zinc nitrite compound?
Dont have zinc to test it out

Colourless prismatic crystals of Zn(NO2)2.2H2O.C6H12N4 are obtained by adding sodium nitrite to the solution of a zinc salt containing hexamethylenetetramine. Zinc nitrite also occurs in the hygroscopic yellow crystalline double salt, 3KNO2.Zn(NO2)2.3H2O, obtained by acting with nitrous acid on zinc hydroxide suspended in potassium nitrite solution, and in the similar salt, 2KNO2.Zn(NO2 )2.H2O, prepared by mixing solutions of potassium nitrite and zinc acetate or nitrate. According to Ray, zinc nitrite can only exist pure and uncombined in dilute solutions. This solution is acid, and the nitrous acid set free by hydrolysis decomposes on concentration -

3HNO2 = HNO3+2NO+2H2O.

Zinc nitrate is thus continuously formed, and the residue from evaporation, even when conducted in vacuo, is a basic nitrate.

A dilute solution of zinc nitrite may be prepared by interaction between solutions of zinc sulphate and barium nitrite, but, though the trihydrate, Zn(NO2)2.3H2O, is said to have been obtained from such a solution, only basic or very impure zinc nitrites have usually been obtained.

The monohydrate, Zn(NO2)2.H2O, was said to be obtained in fine needles by treating a mixture of sodium nitrite and magnesium sulphate with alcohol, filtering, and evaporating in vacuo. The dry salt, its aqueous solution, and (more gradually) its alcoholic solution, give off nitrogen oxides and leave a mixture of zinc nitrate and hydroxide.


Aluminum nitrite decomposes to NO2
Aluminum sulfate and sodium nitrite


Possibilities form a urea complex
Formed amine complex
Or other complex


[Edited on 15-7-2018 by symboom]

XeonTheMGPony - 15-7-2018 at 04:58

Why do you do this?????????????????????????????????????????????????????????

j_sum1 - 15-7-2018 at 14:52

Quote: Originally posted by XeonTheMGPony  
Why do you do this?????????????????????????????????????????????????????????
Good point.
@ symboom, Some sense of context would be wonderful. Either a question or a bit of an explanation of what your experiment is aiming to achieve as well as the reaction conditions. (Is it an experiment? or merely a summary of some stuff you looked up?)

As it is, it is difficult for anyone to contribute meaningfully to the threads you start. Which is a shame because it looks like it might be interesting content.

AJKOER - 16-7-2018 at 13:45

Some chemistry may help to explain the role of alcohol noted above. This honor research work at digitalcommons.iwu.edu/cgi/viewcontent.cgi?article=1287&context=jwprc , may help illuminate the role that alcohol could play, which is basically to scavenge photo-initiated (or possibly otherwise) generated hydroxyl radicals.

Some key reactions could include:

NO2- + H2O + hv --> .NO + .O-

.O- + H2O = .OH + OH-

NO2- + H2O + hv --> .NO2 + e-(aq)

and also, without light, a Fenton-type reaction I suspect (see, for example, "Fenton chemistry in biology and medicine*" by Josef Prousek, and related reaction (15) on page 2330):

Fe(ll)/Cu(l) + HONO --> Fe(lll)/Cu(ll) + NO- + .OH

.OH + NO2- = OH- + .NO2

2 .NO + O2 --> 2 .NO2

2 .NO2 + H2O --> HNO2 + HNO3

all suggesting 'a slow formation of nitrate'.

Also, in the event of working with, in particular, anodic transition metals (like Cu, Fe,...), I would argue that this reaction may also be of import in recycling:

Fe(lll)/Cu(ll) + e-(aq) = Fe(ll)/Cu(l)

The following electrochemical reaction with oxygen and H+ is then possible leading to basic salts formation upon exposure to air:

4 Fe(ll)/Cu(l) + O2 + 2 H+ --> 4 Fe(lll)/Cu(ll) + 2 OH- (see https://www.chegg.com/homework-help/questions-and-answers/co... )

which may accounts for the comment relating to air sensitivity with basic salt formation cited in the opening thread.

Note, zinc itself is considered to be a post-transition or quasi transition metal.

Also, alcohol is an better medium over water to promote solvated electron formation. The superoxide radical anion could then be formed from air interaction and producing further products:

O2 + e- = .O2-

And, with nitric oxide interaction:

.O2- + .NO --> ONOO- (see https://en.wikipedia.org/wiki/Peroxynitrite)

the creation of the peroxynitrite ion, which may be converted to nitrate, which could account, in part, for the comment "Exposure of its dilute solution to air causes slow formation of nitrate". Interestingly, the aqueous decay of the peroxynitrite ion is said to be a source of hydroxyl radicals (see https://www.ncbi.nlm.nih.gov/pubmed/9462930 ). Another source provides a suggested path:

H+ + ONOO- = HONO2 --> .OH + .NO2 (see p.1212 at https://watermark.silverchair.com/dkq075.pdf?token=AQECAHi20... )

Also, the reaction between between NO and H2O2 is claimed also to be a source of .OH (see https://www.ncbi.nlm.nih.gov/pubmed/9545532 ). To quote from the abstract:

"The interactions of H2O2 and NO represent a biologically feasible reaction mechanism that can account for OH-induced damage in cellular environments where transition metal ions are unavailable for participation in the superoxide-mediated Fenton reaction. The ability of the NO/H2O2 complex to generate OH independently of iron or other transition metals provides a new focus for studies concerned with the origin of tissue-specific damage caused by oxygen-derived species."

Here is an insightful attempt to derive the reaction from above:

H+ + .O2- + .NO --> H+ + ONOO- = HONO2 --> .OH + .NO2

H+ + .O2- = .HO2 (pKa=4.88)

.HO2 + .HO2 = H2O2 + O2 (not a particularly fast reaction)

Upon substitution of the 1/2 of the above:

1/2 H2O2 + 1/2 O2 + .NO --> .OH + .NO2

But as 1/2 O2 + .NO = .NO2, the above derived reaction appears to imply that the presence of NO/NO2 is catalytic to the self decomposition of H2O2 into hydroxyl radicals apparently, per above, through a NO/H2O2 complex formation/splitting.

Interestingly, as .OH + .NO = HNO2 (reversible in light forming radicals) = H+ + NO2- , there is a connection to acidified nitrite which is paired with H2O2 to increase cytotoxicity in disease treatment protocols (see, for example, http://www.jbc.org/content/271/11/6144.full.pdf ).

[Edited on 17-7-2018 by AJKOER]

AJKOER - 17-7-2018 at 18:27

Quote: Originally posted by AJKOER  
.....

Also, alcohol is an better medium over water to promote solvated electron formation. The superoxide radical anion could then be formed from air interaction and producing further products:

O2 + e- = .O2-

And, with nitric oxide interaction:

.O2- + .NO --> ONOO- (see https://en.wikipedia.org/wiki/Peroxynitrite)

the creation of the peroxynitrite ion, which may be converted to nitrate, which could account, in part, for the comment "Exposure of its dilute solution to air causes slow formation of nitrate". Interestingly, the aqueous decay of the peroxynitrite ion is said to be a source of hydroxyl radicals (see https://www.ncbi.nlm.nih.gov/pubmed/9462930 ). Another source provides a suggested path:

H+ + ONOO- = HONO2 --> .OH + .NO2 (see p.1212 at https://watermark.silverchair.com/dkq075.pdf?token=AQECAHi20... )

Also, the reaction between between NO and H2O2 is claimed also to be a source of .OH (see https://www.ncbi.nlm.nih.gov/pubmed/9545532 ). To quote from the abstract:

"The interactions of H2O2 and NO represent a biologically feasible reaction mechanism that can account for OH-induced damage in cellular environments where transition metal ions are unavailable for participation in the superoxide-mediated Fenton reaction. The ability of the NO/H2O2 complex to generate OH independently of iron or other transition metals provides a new focus for studies concerned with the origin of tissue-specific damage caused by oxygen-derived species."

Here is an insightful attempt to derive the reaction from above:

H+ + .O2- + .NO --> H+ + ONOO- = HONO2 --> .OH + .NO2

H+ + .O2- = .HO2 (pKa=4.88)

.HO2 + .HO2 = H2O2 + O2 (not a particularly fast reaction)

Upon substitution of the 1/2 of the above:

1/2 H2O2 + 1/2 O2 + .NO --> .OH + .NO2

But as 1/2 O2 + .NO = .NO2, the above derived reaction appears to imply that the presence of NO/NO2 is catalytic to the self decomposition of H2O2 into hydroxyl radicals apparently, per above, through a NO/H2O2 complex formation/splitting.

Interestingly, as .OH + .NO = HNO2 (reversible in light forming radicals) = H+ + NO2- , there is a connection to acidified nitrite which is paired with H2O2 to increase cytotoxicity in disease treatment protocols (see, for example, http://www.jbc.org/content/271/11/6144.full.pdf ).

[Edited on 17-7-2018 by AJKOER]


Here is another interesting attempt to derive the H2O2-NO reaction from above. Starting as before:

H+ + .O2- + .NO --> H+ + ONOO- = HONO2 --> .OH + .NO2

H+ + .O2- = .HO2 (pKa=4.88)

Note, substituting, implies:

.HO2 + .NO --> .OH + .NO2

which is a valid reaction. Now, adding .H to each side, noting that .H + .HO2 = H2O2 and .H + .NO2 = HONO (= H+ + NO2-) produces a potential Scenario ll:

H2O2 + .NO = .OH + HONO

Proceeding, noting that:

.H + .OH = H2O + Photon (see p. 70 at https://books.google.com/books?id=mO6Z07lHQO4C&pg=PA70&a... )

So, alternately, we could have:

H2O2 + .NO = H2O + Photon + .NO2

which suggests aqueous NO2 is a possible answer ignoring the effect of the photon (Scenario lll).

Note, upon choosing to combine .H with .OH (and not .NO2 as above), or equivalently incorporating the photon effect on water, we have:
H2O = H+ + OH-
OH- + Photon → .OH + e-
H+ + e- = .H
-------
Net: H2O + Photon = .OH + .H

which is where we started, so it is logical to combine the .H with .NO2 forming HONO as was first done in Scenario ll.

Now, there are three incongruous results! First, as originally reported .NO2 as well as .NO were suggested as possibly catalytic with an hydroxyl radical product. In Scenario ll, there is no .OH just an aqueous .NO2 radical. Lastly, and there is a hydroxyl radical along with HONO. Unbelievably to me, all of these scenarios are mentioned in the report of research on the H2O2-NO reaction (link: https://www.researchgate.net/publication/23855494_Reaction_o... ) to quote:

“The H2O2-NO reaction was complex. There was an induction period followed by a marked acceleration in reactant removal. The final products of the reaction, NO2, probably H2O, and possibly HONO2, were produced mainly after all the H2O2 was removed. The overall stoichiometry in the presence of excess NO was H2O2 + NO → H2O + NO2. The initial induction period gives an upper limit to the homogeneous gas phase reaction coefficient of 5 × 10-20 cm3/molecule sec for the reaction NO + H2O2 → HONO + HO. The HO radical presumably is removed via HO + NO → HONO. The HONO intermediate was shown to disproportionate to NO2 + NO + H2O in a relatively slow first-order reaction. The acceleration in H2O2 removal after the NO-H2O2 reaction is started is caused by NO2 catalysis: NO2 + H2O2 → HONO2 + HO and HONO2 + NO → HONO + NO2.”

Source: Reaction of hydrogen peroxide with nitrogen dioxide and nitric oxide, an article in The Journal of Physical Chemistry 76(14) · August 1972, DOI: 10.1021/j100658a001,
Available at : https://www.researchgate.net/publication/23855494_Reaction_o... .

Now, my comment above on 'incongruous' scenarios should perhaps be amended to 'seemingly incongruous' as for example, per the comment "the dynamics of OH formation from two photon absorbed NO2 with H2O" at https://pubs.acs.org/doi/abs/10.1021/jp3029825 , there appears to be governance based on the favored radical path under specified conditions. For example, one of my other product results, namely aqueous .NO2, is favored with an excess of NO and another leading to .OH and HONO is promoted in a homogeneous gas phase reaction. My original work suggesting a role for 1/2 O2 + NO, or a NO2 presence, is verified by the article.

[Edited on 18-7-2018 by AJKOER]

AJKOER - 18-7-2018 at 09:20

OK, my suggested path summary as to why transition metal nitrites are unstable in air forming nitrate. I start with the photosensitivity of nitrites, which I suspect contributes to the self-decomposition reaction of HNO2 (also listed) :

NO2- (aq) + hv --> .NO + .O-

.O- + H2O = .OH + OH-

NO2- + H2O + hv --> .NO2 + e-(aq)

Other possible reactions:

.NO2 + H2O --> NO3- + 2 H+ + e-(aq)

Fe(lll)/Cu(ll) + e-(aq) = Fe(ll)/Cu(l)

Net Redox: Cu(ll)/Fe(lll) + .NO2 + H2O --> Cu(l)/Fe(ll) + NO3- + 2 H+ (see http://chemequations.com/en/?s=Fe%3A3%2B+%2B+NO2+%2B+H2O+%3D... )

NO2- + .OH = .NO2 + OH-

2 .NO + O2 --> 2 .NO2

2 .NO2 + H2O --> 2 H+ + NO2- + NO3-

Cu(H2O)6]2+ (aq) + H2O (l) = [Cu(H2O)5(OH)]+ (aq) + H3O+ (aq)

H+ + NO2- = HNO2

2 HNO2 --> NO2 + NO + H2O (other than cold dilute solutions)

3 HNO2 --> HNO3 + 2 NO + H2O (warm concentrated solutions, see https://en.wikipedia.org/wiki/Nitrous_acid)

O2 + e-(aq) --> .O2-

.O2- + .NO --> ONOO- (see https://en.wikipedia.org/wiki/Peroxynitrite)

H+ + ONOO- = HONO2 --> .OH + .NO2 (see https://www.ncbi.nlm.nih.gov/pubmed/9462930 )

Finally, a possible electrochemical reaction with oxygen and H+ is then possible leading to basic salts formation upon exposure to air:

4 Fe(ll)/Cu(l) + O2 + 2 H+ --> 4 Fe(lll)/Cu(ll) + 2 OH- (see https://www.chegg.com/homework-help/questions-and-answers/co... ).

[Edited on 18-7-2018 by AJKOER]