byko3y
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Oxygen in the Na2O is bound to two atoms of sodium, so Na2O is a relatively strong reducing agent, while Cl2 is a strong oxidation agent just like O2,
and NaCl is neutral.
HCl + O2 => Cl2 + H2O
reducing agent + oxidation agent => oxidation agent + reducing agent. However, this reaction requires catalyst and/or some drastic conditions.
And this one does not:
Cl2 + H2O <=> HOCl + HCl
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blogfast25
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For the reaction 2 NaCl + ½ O2 === > Na2O + Cl2, the Standard Enthalpy of Reaction is about +
400 kJ/mol (of Na2O). This, reasonably assuming entropic effects are small, strongly suggests that this reaction is strongly endo-energetic
(Gibbs Free Energy change ΔG > 0), so thermodynamically strongly unfavourable. Any corrections for higher temperatures are unlikely to
fundamentally change that.
No amount of nifty catalysis will change that either, as catalysts do not affect ΔG, they only affect Activation Energy. The equilibrium
constant K of a given equilibrium reaction is not affected by a catalyst.
Blowing hot oxygen through molten NaCl might produce some Cl<sub>2</sub> because the removal of a reaction product
‘pulls’ the equilibrium somewhat to the right (Le Chatelier principle).
[Edited on 5-6-2015 by blogfast25]
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DraconicAcid
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Sodium oxide is not a reducing agent.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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byko3y
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Oxygen in the Na2O is reducing agent only for the first sodium atom, just like water and alcohol.
Na2O + Cl2 => NaOCl + NaCl
There's no way O2 (which tends to have 2 valencies) can oxidize the Cl(-1). But something like singlet oxygen, H2O2, ozone might be able.
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Molecular Manipulations
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I like how you proved that sodium oxide is a reducing agent because two sodium atoms must share one oxygen (which has nothing to due with valance
electrons of course!).
The reaction you showed about oxygen oxidizing hydrogen chloride to water and chlorine (which doesn't even happen, as you seem to know) doesn't really
help your case nor does showing that the opposite reaction can occur (well, not quite opposite, hypochlorous acid has chlorine in the 1st oxidation
state).
You should study thermodynamics.
-The manipulator
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halogen
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F. de Lalande and M. Prud'homme showed that a mixture of boric oxide and sodium chloride is decomposed in a stream of dry air or oxygen at a red heat
with the evolution of chlorine.
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Molecular Manipulations
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Quote: Originally posted by byko3y  | Oxygen in the Na2O is reducing agent only for the first sodium atom, just like water and alcohol.
Na2O + Cl2 => NaOCl + NaCl |
Since you apparently can't draw lewis structures, I'll help you out. That reaction maybe does occur, but only chlorine changes oxidation states.
Sodium is still in +1 and oxygen in -2. One chlorine goes from 0 to -1 the other from 0 to +1, nothing else changed.
-The manipulator
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DraconicAcid
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| Quote: Originally posted by byko3y | Oxygen in the Na2O is reducing agent only for the first sodium atom, just like water and alcohol.
Na2O + Cl2 => NaOCl + NaCl |
I'm sorry, but the only polite way I can respond to this is, "What the hell are you talking about?" In the reaction you've given, oxide is not acting
as either a reducing agent or an oxidizing agent- the chlorine is disproportionating (one chlorine gets oxidized, the other gets reduced).
Please remember: "Filtrate" is not a verb.
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byko3y
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Sorry, you are wrong about HCl oxidation http://pubs.acs.org/doi/abs/10.1021/ie00036a014 The reaction is exothermic.
Basic thermodinamics doesn't explain anything, it just deals with a fact about how much energy consumed/released.
When an atom already has one electron drawn to him, then it cannot hold to another one equally strong. This is why in NaOH the O-H is a coavalent
bound, despite the fact that oxygen has electronegativity of 3.5 and H2 - 2.1.
Na+ + -ONa + Cl2 => Na+ + -O* + NaCl (Na+ + Cl-) + Cl* => 2Na+ + -OCl + Cl-
Na+ + -OH + Cl2 => Na+ + -O* + HCl (H+ + Cl-) + Cl* => Na+ + -OCl + H+ + Cl-
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DraconicAcid
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| Quote: Originally posted by byko3y |
When an atom already has one electron drawn to him, then it cannot hold to another one equally strong. This is why in NaOH the O-H is a coavalent
bound, despite the fact that oxygen has electronegativity of 3.5 and H2 - 2.1. |
No, this is because of the small size of the hydrogen atom and the fact that it only has one electron (so that if it were truly ionic, it would just
be a bare proton, with an insane charge-to-radius ratio). The bond in HF is also covalent, despite the larger electronegativity difference.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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blogfast25
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Polite answer: he doesn't know what he's talking about.
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Molecular Manipulations
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Is it just me or does the OP's (who's post he deleted) username look remarkably like something PHD chemist would make?
About your link... That process takes the term 'catalyst' too far IMHO.
Reacting HCl with MnO2 to produce Cl2 and MnCl2, which can react with hot oxygen to make more Cl2 and MnO2. This is nothing like what you implied, and
the fact that it's over-all exothermic means nothing in light of all the reactions involved.
[Edited on 5-6-2015 by Molecular Manipulations]
-The manipulator
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blogfast25
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halogen:
| Quote: | | F. de Lalande and M. Prud'homme showed that a mixture of boric oxide and sodium chloride is decomposed in a stream of dry air or oxygen at a red heat
with the evolution of chlorine. |
Care to provide a reference for that?
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byko3y
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DraconicAcid, I wanted to use terms polar(which includes ionic)-nonpolar instead of coavalent-ionic, because obviuosly you can't judge about whether
the coupound is ionic by electronegativities of its atoms.
>Sodium is still in +1 and oxygen in -2. One chlorine goes from 0 to -1 the other from 0 to +1, nothing else changed
Molecular Manipulations, integral oxidation states are too primitive to describe what actually happens.
In the NaOCl molecule chlorine is close to zero oxidation state. To convert the compounds back to chlorine, you need to oxidize them with H+,
substituting an electron donating sodium with not so electron donating hydrogen.
To verify the theory let's take reaction 2NaOCl => 2NaCl + O2, supposing Na(+1), O(-1), Cl(0) in the NaOCl
Na+ + -OCl => Na+ + -Cl + O
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DraconicAcid
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| Quote: Originally posted by byko3y |
In the NaOCl molecule chlorine is close to zero oxidation state. To convert the compounds back to chlorine, you need to oxidize them with H+,
substituting an electron donating sodium with not so electron donating hydrogen. |
In NaOCl, the sodium ion is not particularly associated with any hypochlorite anion. It is not a molecule. The sodium is not being "electron
donating".
Adding H+ does not oxidize the chlorine, it allows the chloride ion to react with the hypochlorite ion to form water and chlorine.
| Quote: | To verify the theory let's take reaction 2NaOCl => 2NaCl + O2, supposing Na(+1), O(-1), Cl(0) in the NaOCl
Na+ + -OCl => Na+ + -Cl + O |
You're not verifying anything here, other than you don't know what you're talking about.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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blogfast25
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| Quote: Originally posted by byko3y |
To verify the theory let's take reaction 2NaOCl => 2NaCl + O2, supposing Na(+1), O(-1), Cl(0) in the NaOCl
Na+ + -OCl => Na+ + -Cl + O |
How can Cl be Cl(0) [in ClO<sup>-</sup>] when it's bound to a more electronegative atom (oxygen)?
NaClO === > NaCl + 1/2 O<sub>2</sub>
... is more exo-energetic (Gibbs) than:
NaClO === > 1/2 Na<sub>2</sub>O + 1/4 O<sub>2</sub> + 1/2 Cl<sub>2</sub>
Which explains the outcome, albeit not the mechanism.
[Edited on 5-6-2015 by blogfast25]
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blogfast25
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| Quote: Originally posted by Molecular Manipulations |
About your link... That process takes the term 'catalyst' too far IMHO.
Reacting HCl with MnO2 to produce Cl2 and MnCl2, which can react with hot oxygen to make more Cl2 and MnO2. |
This is indeed not catalysis. He's already shown he doesn't really understand how a catalyst works by claiming a thermodynamically unfavourable
reaction can be made to work anyway by means of a 'byko-catalyst'! 
[Edited on 5-6-2015 by blogfast25]
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byko3y
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blogfast25, http://pubs.acs.org/doi/abs/10.1021/ie00036a014
| Quote: | | How can Cl be Cl(0) [in ClO-] when it's bound to a more electronegative atom (oxygen)? |
That's what I'm talking about, the oxygen already has an excessive electron, he has a different electronegativity now, he has been reduced, he is not
so much electronegative anymore. https://commons.wikimedia.org/wiki/File:Sodium-hypochlorite....
| Quote: | | The sodium is not being "electron donating". |
Then what does the sodium doing there? He's giving off
electrons to Cl- and OCl-.
| Quote: | | it (H+) allows the chloride ion to react with the hypochlorite ion |
How? Why?
I was not telling that H+ oxidizes only chlorine, I was talking about the whole mixture. My theory works perfectly: you oxidize the mixture, leading
to a much more oxidized mixture with overoxidized HOCl with H(+0.7), O(-1), Cl(0) and slightly oxidized (and ionic) HCl with H(+1), Cl(-0.7). (number
are not precise and pretty much relative to the reaction)
In this situation the HOCl is attacked by Cl-, not by H+ (the latter scenario leads to O2 release, and the proton cannot attack anything):
HOCl + Cl- + H+ => HOClCl- + H+ => Cl2 + HO- + H+
Why exactly HOClCl- => Cl2 + HO-? Well, there's probably an equilibrium, but it is strong shifted towards forward reaction (Cl2 production),
because chlorine is unable to abstract electron from the unreduced oxygen (leading to HO*), so of four atoms the oxygen is the one that gets the
negative charge, while ClCl is left without a charge.
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halogen
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Covalent and ionic bonding is not a strict demarcation; while I had understood the reason for this as the instability of the Na20 crystal as opposed
to salt, what byko says is reasonable to me. Inductive effects are well known.
The quote, I am almost sure, was from a book in scimad library (thanks!!), probably "A Text-Book of Inorganic Chemistry", but I haven't found it on
quick look through the chlorine section of the book on halgens, where I expected it. Nevertheless, I didn't make it up.
But! From p. 109 some intresting stuff.
| Quote: | It is not always necessary to use the free metal ; for chlorine also attacks
various compounds, producing chlorides. For example, very often the
metallic oxides are decomposed when heated in chlorine gas.
Silver oxide, Ag20, at ordinary temperatures, is converted into chloride
with liberation of oxygen, while the oxides of the alkaline earths, CaO, BaO,
SrO, are decomposed on heating, with the formation of the corresponding
chlorides and the liberation of oxygen gas, whose volume amounts to half the
volume of the absorbed chlorine.
CaO, BaO, SrO, when heated red-hot and plunged into chlorine, are
actually decomposed with the production of incandescence. MgO is also
decomposed when highly heated in chlorine gas, but without incandescence
appearing. Also the oxides of zinc, cadmium, lead, copper, and nickel are
easily decomposed by hot chlorine, while cobalt and iron oxides are
decomposed only with difficulty. Aluminium oxide and silica are not
decomposed, even at a white heat only small amounts of chlorides being
formed. The trioxides Cr2 8, Mo03, W03 form oxychlorides. SnO, Sb2 5,
As2 3 , when strongly ignited in .the gas, are completely converted into
chlorides.
Several oxides which are only slightly attacked in the presence of hot
chlorine are readily decomposed if mixed with charcoal, and then the mixture
is ignited in a stream of chlorine. For example, SiCl4, SnCl4, A1C13, ThCl4,
etc., may be prepared by this method from the corresponding chlorides.
Chlorine also acts on most bromides and iodides of metals, displacing
bromine or iodine and producing chlorides. Many sulphides are decomposed
when heated in a stream of chlorine, and this provides a method of producing
chlorides of many metals.
A modification of this method consists in passing a stream of chlorine
charged with vapours of disulphur dichloride over heated oxides and oxygen
salts. Many anhydrous chlorides have been obtained in this manner. Si02,
for example, at a dull red heat yields SiCl4 ; A12 3 gives A1C13, while Th02
at a red heat yields ThCl4. W03 gives W02C1 2 at a high temperature, and
WOC14 at a lower temperature. NiO, CoO at 400 give the chlorides, while
Cr2 3 and Fe2Cl3 give CrCl3 and FeCl3 even below redness.
Similarly the oxy-salts of calcium, strontium, and barium (e.g. the
sulphates), are transformed into chlorides by this method. In the application
of this, however, care must be taken not to have excess of disulphur
dichloride if the chlorides to be prepared are volatile. When, however,
the chlorides to be prepared are fairly non-volatile, the action is accelerated
by excess of the sulphur chloride. |
And...
p. 112
| Quote: | Many chlorides which appear stable enough when heated in an inactive
gas decompose readily enough when heated in dry oxygen gas. Thus aluminium
chloride at a dark red heat in oxygen gas evolves part of its
chlorine, while such a stable chloride as that of sodium when heated very
intensely in dry air also evolves chlorine. (1) |
1. de Sanderval, Compt. rend., 1893, Il6, 641.
[Edited on 5-6-2015 by halogen]
F. de Lalande and M. Prud'homme showed that a mixture of boric oxide and sodium chloride is decomposed in a stream of dry air or oxygen at a red heat
with the evolution of chlorine.
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blogfast25
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halogen:
Why do you believe the Na2O crystal is substantially different from NaCl? Even their Enthalpies of Formation are similar (mainly lattice energy). Not
to say they are the same of course...
| Quote: | Several oxides which are only slightly attacked in the presence of hot
chlorine are readily decomposed if mixed with charcoal, and then the mixture
is ignited in a stream of chlorine. For example, SiCl4, SnCl4, A1C13, ThCl4,
etc., may be prepared by this method from the corresponding chlorides. |
This is an entirely different type of reaction, as well you know.
Much of the rest of the quote has NOTHING to do with the OP's question. Less is more.
Of course oxygen can displace chlorine in some cases, or vice versa: where the Gibbs Free Energy change is negative it's usually possible, assuming no
kinetic obstacles. NaCl to Na2O just doesn't appear to be one of them.
You sport a motto you can't even find the reference for? Nought queerer than folk! 
[Edited on 5-6-2015 by blogfast25]
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byko3y
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blogfast25, Gibbs free energy changes with change of temperature and pressure. This is why water vaporizes on heating and condenses on cooling.
The main problem with OP post was that he assumed the reaction needs to be 2NaCl + O2 => Na2O + Cl2, while in fact it is 2NaCl + O2 <=> NaOCl
+ NaCl <=> 1/3 NaClO3 + NaCl, however, at normal conditions this reaction is not observed.
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blogfast25
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I already indicated that in my first post. Learn to read. For the hypothetical reaction the OP enquired about, Gibbs Free Energy change can be
considered relatively invariant of temperature (but low long is a piece of string, of course). Lower pressure would slightly favour it.
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byko3y
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Oxygen and chlorine are gases with different molar mass, while NaCl and NaOCl are solids. This way we have different changes in pV. And entropy is
defined by a hell lot of factors, including solvent.
But of course OP ment something like room temperature and atmospheric pressure, and in those conditions energies were explicitly measured, and of
course you can't convert NaCl to Na2O, because you can't.
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halogen
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I just assumed they didn't fit together well.
F. de Lalande and M. Prud'homme showed that a mixture of boric oxide and sodium chloride is decomposed in a stream of dry air or oxygen at a red heat
with the evolution of chlorine.
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DraconicAcid
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| Quote: Originally posted by byko3y |
| Quote: | | The sodium is not being "electron donating". |
Then what does the sodium doing there? He's giving off
electrons to Cl- and OCl-. |
No, it's just sitting there, being a positive ion, attracting the hypochlorite enough that you have a crystalline solid with no overall charge.
| Quote: | | Quote: | | it (H+) allows the chloride ion to react with the hypochlorite ion |
How? Why?
I was not telling that H+ oxidizes only chlorine, I was talking about the whole mixture. My theory works perfectly: you oxidize the mixture, leading
to a much more oxidized mixture with overoxidized HOCl with H(+0.7), O(-1), Cl(0) and slightly oxidized (and ionic) HCl with H(+1), Cl(-0.7). (number
are not precise and pretty much relative to the reaction) |
Your "theory" is that you make up some numbers and spew some bafflegab. You really have no idea what you are talking about, and I wash my hands of
the thread.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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