December 05, 2020, 09:37:48 AM
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Topic: Need help identifying strange side-reaction happening in my chemistry practical  (Read 920 times)

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Offline avanti262

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Hiii! Just to provide some context, i am currently studying as a Y5 under the IB program and am doing my chemistry IA. My experiment is to determine the rate of catalysed decomposition of hydrogen peroxide (H2O2).

The experiment is simple; in a closed chamber, first drop in the catalyst and turn on the stir-bar to ensure mixing is constant. Then insert the H2O2 diluted with di-water with a syringe to avoid having any time lag. Observe the change in pressure. The two catalyst that I have chosen to study are Iron(II) Oxide and Titanium(IV) Oxide.

However, while carrying out the experiment, I have encountered a strange interaction when H2O2 to above 65⁰c where oxygen gas is evolved and then quickly consumed by some unknown side reaction. The normal pressure of my chamber would begin around 101 kPa, followed by an increase to around 1010-115 kPa when the catalyst is introduced, and then quickly, the pressure in the chamber drops to around 90 kPa. It is interesting to note that this only occurs above 65⁰c.

I have ensured that my reaction chamber has no leakages with vacuum grease. The solution is constantly mixed with a stir-bar spinning at 300 rpm. The chamber is heated by a water bath that is heated over a heating mantle, the temperature is monitored via temperature probe.

This interaction is present consistently across multiple runs. I have done thorough research on the mechanism behind the catalysed decomposition of H2O2 using the two mentioned catalysts. None of the papers that I have read through mentioned this problem.

Any ideas as to what is going on here?

Offline chenbeier

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Titandioxide can react wirh peroxide.


https://pubs.acs.org/doi/10.1021/j100585a011

Ti(O2)2+ can be formed

Offline avanti262

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I see, but how would this account for the loss in pressure of the chamber. It also doesn't explain in the Iron(II) Oxide runs how the pressure in the chamber had also decreased in similar fashion.

But thank you very much for you're *delete me* Really appreciate it :))!

Offline Enthalpy

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Maybe you could evaluate the chamber volume, the amount of oxygen expected, and the amount corresponding to the pressure changes. That would tell you what kind of explanations are possible or not.

Offline billnotgatez

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Is there a change in color?

It may not be important but here is a WIKI quote

https://en.wikipedia.org/wiki/Iron_oxide#:~:text=Iron%20oxides%20are%20chemical%20compounds,many%20geological%20and%20biological%20processes.
Quote
Iron oxides are chemical compounds composed of iron and oxygen. There are sixteen known iron oxides and oxyhydroxides, the best known of which is rust, a form of iron(III) oxide.

Offline avanti262

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Maybe you could evaluate the chamber volume, the amount of oxygen expected, and the amount corresponding to the pressure changes. That would tell you what kind of explanations are possible or not.

Im certain that the pressure of the chamber should increase as the reaction moves forward and remain at a higher pressure than before. According to my calculations, a complete decomposition of 5ml 6% H2O2 will yield around 100cm^3 of O2. About the volume of the chamber, I am not very sure and will have to go back to the lab to take a look at it.

I do know for a fact that from 35C to 55C, the runs follows the expected trend of the pressure increasing at increasing rates as the temperature increases and holding the pressure just below the maximum evolved. However, at 65C, the pressure will first increase then decrease. This is illustrated below by the graph of the experiment for Titanium(IV) Oxide at 72C:

https://drive.google.com/open?id=19ceFpsq0xsPe7pJl50R8TtkYQS7Qh1bj
*sorry for the linking, dont understand how to attatch an image.

It must be noted the the original atmospheric pressure was 101 kPa. The experiment started at a slightly higher pressure due to the fact that some Hydrogen Peroxide was already decomposing from the high temperature.
« Last Edit: October 15, 2020, 04:41:51 AM by avanti262 »

Offline avanti262

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Is there a change in color?

It may not be important but here is a WIKI quote

https://en.wikipedia.org/wiki/Iron_oxide#:~:text=Iron%20oxides%20are%20chemical%20compounds,many%20geological%20and%20biological%20processes.
Quote
Iron oxides are chemical compounds composed of iron and oxygen. There are sixteen known iron oxides and oxyhydroxides, the best known of which is rust, a form of iron(III) oxide.

Unfortunately I did not observe any sort of color change during the reaction. The solution just turned into whatever color the catalyst was. For Iron(II) Oxide it turned reddish brown and for Titanium(IV) Oxide it turned white. Throughout the entire reaction there was no other color change.

Thank you for the suggestion though!

Offline Borek

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For Iron(II) Oxide it turned reddish brown

That suggests oxidation of Fe(II) to Fe(III) - hardly surprising in the presence of water and oxygen.

No idea about the Ti part.
ChemBuddy chemical calculators - stoichiometry, pH, concentration, buffer preparation, titrations.info, pH-meter.info

Offline Enthalpy

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But this happens with TiO2 too, from which I don't expect the same gag as from iron oxides.

Also, oxygen evolves from the liquid, and I suppose the catalyst is immersed. Though, the pressure first increases, then drops, despite the catalyst looks inaccessible to the evolved oxygen.

Could something else, outside the liquid, scavenge the oxygen? A seal ring, a grease...?

Does the temperature drop between 115 kPa and 90 kPa, and enough to explain the pressure drop? Then, just a leak appearing only at 115 kPa and +65°C would suffice.

Offline billnotgatez

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@avanti262
I am trying to understand some points of the experiment
how big is the chamber
how much iron oxide did you use
is the iron oxide some crystals or a powder or water slurry
is there air in the head-space of the chamber
You used 5ml 6% H2O2
did the color of the iron oxide change from sort of black to sort of red
how did you apply heat to the chamber ( was it an electric)
if you have a picture it would be nice to see it

Offline avanti262

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@avanti262
I am trying to understand some points of the experiment
how big is the chamber
how much iron oxide did you use
is the iron oxide some crystals or a powder or water slurry
is there air in the head-space of the chamber
You used 5ml 6% H2O2
did the color of the iron oxide change from sort of black to sort of red
how did you apply heat to the chamber ( was it an electric)
if you have a picture it would be nice to see it

Hello,

I will answer your questions from top to bottom:

The chamber was a two-necked 100ml round bottomed flask.

The amount of iron oxide and titanium oxide used in each run was 0.200g and 0.500g respectively.

The iron oxide and titanium oxide crystals were in powder formed, then submerged fully by the diluted hydrogen peroxide.

Yes, the chamber was not vacuumed and had atmospheric air from the surroundings inside.

Yes, it was 5ml of 6% hydrogen peroxide dilluted with 15ml of di-water

Yes, the iron oxide had a slight color change from red to blackish red. However, titanium oxide remained white.

The heat was applied to the chamber via water bath that was heated by a heating mantle. Temperature was monitored via thermometric probe.

Here is a picture of the setup:https://drive.google.com/file/d/1Gpk1iPyFyjAd9rPdSrdI5W2UipBwOXbd/view?usp=sharing

And here is a more detailed drawn diagram of the above setup: https://drive.google.com/file/d/1ZwQpW4tu1FK1NEoqY8lYrC5RUnxpsa8l/view?usp=sharing
« Last Edit: October 21, 2020, 01:17:53 AM by avanti262 »

Offline avanti262

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But this happens with TiO2 too, from which I don't expect the same gag as from iron oxides.

Also, oxygen evolves from the liquid, and I suppose the catalyst is immersed. Though, the pressure first increases, then drops, despite the catalyst looks inaccessible to the evolved oxygen.

Could something else, outside the liquid, scavenge the oxygen? A seal ring, a grease...?

Does the temperature drop between 115 kPa and 90 kPa, and enough to explain the pressure drop? Then, just a leak appearing only at 115 kPa and +65°C would suffice.

That was my first thought at first, but I thought that it seemed a bit strange for a leakage to only occur at temperatures above 65°C. I had ensured that each fitting in the setup was sealed using vacuum grease which should help with leakage issues.

Take note that during extended trails (+1min), the chamber was able to withstand and hold a pressure of 120 kPa and that the temperature was maintained at each specific point.

Offline billnotgatez

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@avanti262
Thank you for the very nice presentation of your project.

By the way do know how many moles of O2 you expect and how many moles of Iron (II) oxide you used.

Offline avanti262

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By the way do know how many moles of O2 you expect and how many moles of Iron (II) oxide you used.

All calculations made here are purely theoretical, they may very as Hydrogen Peroxide is known to undergo photo-decomposition. I have tried to minimise this by storing the solution in a dark bottle during experiments.

The amount of moles for the Hydrogen Peroxide solution (5ml diluted with 15ml of water) is 0.00881 (3sf).

The concentration of the solution 0.441 mol/dm3 (3sf)

The moles of oxygen evolve from the decomposition is 0.00441

The moles of the Iron(II) Oxide would be 0.00278 (3sf)


Offline billnotgatez

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@avanti262
I am following this thread and wonder if you have any more thoughts.


Just a side note (sanity check)
 and if I did the math correctly using Ideal gas and nothing else
After rearranging the variables of the Ideal gas law
 (P2) = (P1 . V1 . n2 .T2) / (V2 . n1 . T1)
assuming the container size and moles stay the same then
 (P2) = (P1 .T2) / (T1)
assuming start is 20 C (293 K)and 101 kPa with end being  65 C (338 K)
(P2) = (101 x 338) / (293)
(P2) = about 116.5 kPa
« Last Edit: October 28, 2020, 03:21:21 AM by billnotgatez »

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