# Demonstrating decomposition of hydrogen peroxide using iron(III) nitrate catalyst

I need a way to prove/show that hydrogen peroxide was decomposed through use of catalyst.

I want to ensure that my catalyst: $\ce{Fe(NO3)3}$ or iron(III) nitrate is a catalyst, not a reactant/ consumed during the reaction.

$$\ce{2 H2O2 (aq) ->[Fe(NO3)3 (s)] 2 H2O (l) + O2 (g)}$$

• When the reaction is happening, I will introduce a wooden glowing splint over the bubbling reaction and the wooden splint glowing brighter or reigniting will show that oxygen is being produced. Question:

• Now, how do I show that water is produced? Would I just boil the product I get after reaction (image above) and put cobalt chloride paper at the water vapour?

• Also, if I were to use the orange solution above again as a catalyst (since it still contains iron(III) nitrate), would hydrogen peroxide decompose again? If so, is there a way to put iron(III) nitrate back to its solid state? or any way to reuse as catalyst?

Edit:

If it is easier to answer, it doesn't have to be with iron(III) nitrate. I have an option to use manganese dioxide, which is another catalyst that I can substitute for iron(III) nitrate. I think it should be OK since it does the exact same reaction.

Detection of oxygen: Detection of $\ce{O2}$ by a glowing splint is a good way to detect the oxygen. Also you could capture the gas by a simple fixture e.g. and demonstrate the volume change in the receiver. This way you can actually measure moles $\ce{O2}$ produced (by $PV=nRT$) then moles $\ce{H2O2}$ decomposed stoichiometrically by the formula you've written.

Choice of catalyst: Using $\ce{MnO2}$ would be better if you want to make sure you have a catalyst. $\ce{MnO2}$ will not be consumed during the decomposition; I'm not sure about iron(III) nitrate.

Or, detecting change in $\ce{Fe(NO3)3}$ concentration: Addition of a very small concentration of potassium thiocyanate, $\ce{KSCN}$ (say 1/100 of your $\ce{Fe^{+3}}$ concentration) will yield a deep red product, iron thiocyanate ($\ce{Fe(SCN)^{+2}}$). If you have access to a spectrometer you can measure absorbance of the initial solution's product and the product after decomposing $\ce{H2O2}$.

Prove water was produced: The best way I can think of is to measure the (subtle) change of density of your solution before and after decomposition. Using $\ce{MnO2}$ would make this easy because you can remove/filter it as a solid after decomposition and thus measure mass/volume of your solution before and after. $\ce{H2O2}$ and $\ce{H2O}$ have small albeit detectably different densities at RT. $\ce{H2O2}$ is more dense than water so your density should decrease.

• How would i remove/filter manganese dioxide as solid from water? Is there any specific equipment? Or wcan i just boil it to remove water? – didgocks Jan 25 '17 at 12:22
• Just filter paper through a funnel, maybe a vacuum if you want to be thorough. – khaverim Jan 25 '17 at 16:02
• Would there be a way to prove water for iron(III) nitrate? – didgocks Jan 25 '17 at 20:50
• You can probably still detect a small decrease in density even with iron(III) nitrate in solution – khaverim Jan 25 '17 at 20:54

Yes, apart from potassium iodide which is commonly used as a catalyst in the decomposition of hydrogen peroxide, $\ce{Fe^3+}$ salts, manganese dioxide and nickel hydroxide can be used as a catalysts as alternatives. Since, iron nitrate contains $\ce{Fe^3+}$, it can be used as catalyst.

There are some papers that discuss the use of $\ce{Fe^3+}$ salts as catalyst. Following is the relevant information from the papers:

We should consider the role of the Ferric Chloride ($\ce{FeCl3}$) as catalyst in the decomposition reaction of hydrogen peroxide.(...) The fact is that Iron can exist in two different oxidation states, $\ce{Fe^2+}$ (Ferrous) and $\ce{Fe^3+}$ (Ferric), allows the catalyst to break the reaction into two different redox steps, each of which has a lower energy barrier to completion than the uncatalyzed reaction:

$$\ce{H2O2(aq) + 2Fe^3+(aq) -> O2(g) + 2 Fe^2+(aq) + 2H+(aq)}$$ $$\ce{H2O2(aq) + 2 Fe^2+(aq) + 2 H+ (aq) -> 2H2(l) + 2Fe^3+(aq)}$$

Note the first step in the catalyzed reaction involves reduction of the Ferric Ion ($\ce{Fe^3+}$) to the Ferrous Ion ($\ce{Fe^2+}$), which is then re-oxidized to Ferric Ion ($\ce{Fe^3+}$) in the second step. Hence, on net, the catalyst is not consumed during the course of the decomposition.

$\ce{Fe^3+}$ ions is actually a homogeneous catalyst. The catalytic decomposition of hydrogen peroxide can be essentially explained by two different mechanisms based on the mutual redox transition Fe(III)/Fe(V) (KREMER-STEIN mechanism) and Fe(III)/Fe(II) (HABER-WEISS mechanism), respectively.

According to the mechanism proposed by KREMER and STEIN an intermediate oxygen complex of iron with oxidation number +V is primarily formed by the reaction of $\ce{Fe^3+}$ with $\ce{H2O2}$. This complex reacts with another $\ce{H2O2}$ molecule to water and oxygen thereby reforming $\ce{Fe^3+}$.

$$\ce{Fe^3+ + H2O2 <=> [Fe^{III}OOH]^2+ + 2H+ <=> [Fe^{V}O]^3+ + H2O ->[H2O2] Fe^3+ + 2H2O + O2}$$

According to the mechanism proposed by HABER and WEISS the $\ce{Fe^3+}$ ions initiate a radical reaction, after which the chain reaction consumes the hydrogen peroxide. This mechanism can explain the high reaction rate very well.

Chain initiation: $\ce{Fe^3+ + H2O2 <=> [Fe^{III}OOH]^2+ + 2H+ <=> Fe^2+ + HOO. + H+}$

Chain propagation: $\ce{Fe^2+ + H2O2 -> Fe^3+ + 2OH.}$ $\ce{Fe^3+ + H2O2 + OH. -> Fe^3+ + HOO. + H2O -> Fe^2+ + H+ + O2 + H2O }$