chemical production of ozone

 
ozone, when liquified at -112 °C, is an indigo-violet liquid which is dangerously explosive.
ozone gas is soluble in carbon tetrachloride.
 
For more information about the stability of liquid ozone see: https://sites.google.com/site/energeticscribble/rocket-oxidizers. For some information about ozonides, see https://sites.google.com/site/energeticchemical/ozonide-superoxide
 
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  28 gms. of ozone per cubic metre of oxygen evolved,
by gently heating small quantities of powdered barium per-
oxide in eight times its volume of concentrated sulphuric
acid. Hydrogen peroxide is likewise formed in small quan-
tities under these conditions:—
4H2SO4 + 4BaO2 -> (i) 4BaS04 + 4H2 O + 2Oa
(ii) 4BaSO4 + 4H2 O + 0 3 + 0
(iii) 4BaSO4 4- 4 H2 O2 .
The same investigator showed that similar results wore
obtained with other peroxides, notably those of magnesium,
zinc, sodium, and potassium.
Even better results can be obtained by the decomposition
and gentle dehydration of permanganic acid or potassium
dichromate,
Mn2O7 -> 2MnO2 + O3.
by gently heating small quantities of powdered barium peroxide in eight times its volume of concentrated sulphuric acid. Hydrogen peroxide is likewise formed in small quantities under these conditions
4H2SO4 + 4BaO2 -> (i) 4BaS04 + 4H2 O + 2Oa
 4BaSO4 + 4H2 O + 0 3 + 0
4BaSO4 4- 4 H2 O2
Similar results were obtained with other peroxides, notably those of magnesium, zinc, sodium, and potassium. Even better results can be obtained by the decomposition and gentle dehydration of permanganic acid or potassium dichromate.
Mn2O7 -> 2MnO2 + O3
As dehydrating agent, sulphuric acid is niost conveniently employed in the proportions of one of potassium permanganate to two of sulphuric acid. De la Coux (" L'Ozone," p67) states that oxalic acid can be likewise employed in the proportion of 10 gms. of permanganate to 15 gms. of oxalic acid, and that 90 c.c. of oxygen containing 3 mgm. of ozone can be obtained from this mixture
Satisfactory yields of ozone may also be obtained by cautious addition of barium peroxide to a solution of potassium permanganate in sulphuric acid, of density 1*85
 
From the reaction of fluorine and water  An ozone content of upwards of 14 per cent, can be obtained if temp maintained at 0degC.
 
By the thermal decomposition of the persulphates, small
Twenty gms. of dry and freshly prepared ammonium persul
Satisfactory yields of ozone may also be obtained by the
cautious addition of barium peroxide to a solution of potas-
sium permanganate in sulphuric acid, of density 1*85.
By the thermal decomposition of the persulphates, small
quantities of ozone are likewise disengaged, Malaquin (" J.
Pharm. Chem.," VII, 3, 329, 1911) gives the following details
for the preparation of ozonised oxygen by this means.
Twenty gms. of dry and freshly prepared ammonium persul-
phate are mixed with 15 gms. of nitric acid in a small flask;
the air is subsequently displaced by carbon dioxide, and the
mixture cautiously raised to 65° to 70° C. The reaction,
which is strongly exothermic, proceeds somewhat vigorously
when once started, and the resulting oxygen, after removal
of the carbon dioxide, contains 3 to 5 per cent, of ozone and
small quantities of nitrogen.
'
Moissan, in his researches on the properties of fluorine,
which he isolated by the electrolysis of fused potassium
hydrogen fluoride, noted that appreciable quantities of ozone
were produced when a few drops of water were introduced
into an atmosphere of fluorine.
The formation of ozone proceeding according to the
equation:—
3F2 + 3H2O - 6HF + O8,
 
is especially marked at low temperatures, when the rate of
thermal decomposition of any ozone formed is considerably
reduced.
An ozone content of upwards of 14 per cent, in the oxy-
gen disengaged by means of this reaction may be obtained,
if the temperature be maintained at 0° C. De la Coux
 
 
Oxidizing Power of Alkaline Hydrogen Peroxide
 
Although alkaline peroxides (such as CaO2) are stable in the absence of water, hydogen peroxide slowly decomposes in aqueous alkaline solution. A mixture of hydrogen peroxide and ammonium hydroxide (in a 1:3 ratio) acts as a reactive oxidizer, which can attack organic compounds and elemental carbon. The reaction rate is negligible at room temperature, but when heated to 60°C the reaction becomes vigorous and self-sustaining. Such solutions are sometimes known as "base piranha".
With a in 1:1:5 volume ratio of NH4OH, H2O2, and H2O, respectively, the half-life times of peroxide were 4 hours at 50°C and 40 minutes at 80°C. "Reaction of Ozone and H2O2 in NH4OH Solutions and Their Reaction with Silicon Wafers"
Japanese Journal Applied Physics. 43 (2004) pp. 3335-3339 
Magnesium hydroxide inhibits the formation or reactive radicals in alkaline solutions of hydrogen peroxide, interrupting the free radical chain reactions by catching the superoxide anion radicals. Zeronian SH & Inglesby MK (1995) "Bleaching of cellulose by hydrogen peroxide". Cellulose 2: 265-272.
 
Some possible reactions are:
HO2[-] + H2O2 <==> H2O3 + OH[-] 
H2O3 <==> HO3[-] + H[+]
H2O3 + H2O <==> (2)H2O + O21
HO3[-] --> OH[-] + O21
HO2[-] + H2O2 --> OH[-] + O21
HO2[-] + HO3[-] <==> (2)O2[-] + H2
Note that O21 denotes the excited singlet state of the diatomic oxygen molecule, the excited state has an extremely long half-life of 72 minutes in the gas phase, although it quickly decays in solvents, and conveys higher chemical reactivity to the molecule.
 
Elemental titanium shows good corrosion resistance against alkaline hypochlorite. Interestingly, however, alkaline solutions of hydrogen peroxide are known to corrode titanium.  It is thought that the HOO- anion forms a complex with the protective surface oxide layer, Ti(OH)2O2 (with titanium still in the +4 oxidation state), the soluble complex later hydrolyzing into Ti(OH)4 and H2O2.
 
 
Reactions of Chlorate with Hydrogen Peroxide

Some research that should completely put to rest any questions about the reaction between potassium perchlorate and hydrogen peroxide:

Chlorates, when heated to moderate temperatures, disproportionate into perchlorate and chloride.
4NaClO4 --> NaCl + 3NaClO4
Fowler and Grant found that on heating chlorate with silver oxide that the chlorate was completely converted to perchlorate without loss of oxygen, metallic silver also forming.
J. Chem. Soc. 57, 272 (year 1890)

A solution containing one gram sodium chlorate, 1 cc sulfuric acid (specific gravity 1.82), and one gram potassium permanganate in 100 cc water was boled for 30 minutes. The solution showed no perchlorate present.

Hydrogen peroxide failed to oxidize chlorate to perchlorate under alkaline, neutral, or acidic conditions, although minute traces of perchlorate did form under acidic conditions.
“Electrolytic Formation of Perchlorate” C. W. Bennett, E.L. Mack. Chemical Engineer, volume 23, p206

Alkaline
A boiling solution of sodium peroxide failed to oxidize chlorate to perchlorate.

A solution containing one gram sodium chlorate and 1cc ammonium hydroxide (specific gravity 0.90) in 15 cc hydrogen peroxide (30%) was boiled for 30 minutes. Analysis of the solution failed o show any traces of perchlorate, thus showing that alkaline hydrogen peroxide is not a sufficiently powerful oxidizer to convert chlorates to perchlorates.
Nuetral
One gram of sodium chlorate was dissolved in 15 cc hydrogen peroxide (30%) and the solution evaporated to dryness on the water bath. The residue was dissolved in a second 15 cc portion of hydrogen peroxide and again brought to dryness. After dissolving in water, and analysis was performed. This experiment showed that chlorate in neutral solutions is not oxidized to perchlorate by 30% hydrogen peroxide.

Acidified
One gram sodium chlorate was dissolved in 25 cc of hydrogen peroxide (30%) which had previously been acidified wih 1 cc sulfuric acid (specific gravity 1.82). The solution was boiled for one hour. Soon after the solution had reached he boiling point a yellow gas was evolved. This was at first thought to be chlorine but more careful examination showed it to be a mixture of chlorine dioxide and chlorine. Analysis showed small traces of perchlorate had formed. This experiment showed that chlorate, through the action of acidic solutions of hydrogen peroxide, is largely converted to chloride. A considerable amount of chlorine and chlorine dioxide is evolved at the same time. Acidic solutions of 3% hydrogen peroxide also were shown to reduce chlorate to chloride.

It has been suggested that in the above reaction the intermediate formation of small amounts of hydrogen chloride interferes with the reaction, catalytically causing decomposition of the chlorate. Chlorine is known to react with hydrogen peroxide to form hydrochloric acid and oxygen gas. The hydrochloric acid thus formed would attack the remaining chlorate, the products of the reaction being chlorine and chlorine dioxide, the chlorine then reacting with more hydrogen peroxide to again form hydrogen chloride.

Reaction of concentrated sulfuric acid with sodium chlorate did not produce any perchlorate, but it has been reported by other sources that perchlorate is indeed produced. This may be due to different acid concentrations and ratios of reactants.

The reaction between solutions of chloric acid (HClO3) and hydrogen peroxide does not have any appreciable reaction rate until a temperatures above 70degC. (note that perchlorate is not a reaction product in the decomposition reaction, although it may be likely that traces are formed) experiment by Sand. Zelt phys. Chem.,50, 465 (year 1904)

 
 
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