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MgSO4 as an oxidizer?


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Depends on what you mean by ;oxidizer'.
Oxidizer potential is the affinity of an atom/molecle for the electrons of a 'reducer' atom/molecule, so as to fill its outer valence band.

Oxidizer/reducer is a sliding scale.
But if the outer valence orbital of MgSO4 is full, even containing oxygen, it will not readily undergo red-ox reactions.

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10 hours ago, megfish21 said:

I was wondering if magnesium sulphate can be used as an oxidizer in fuel-oxidizing agent mixtures since it contains oxygen.

There are micro-organisms that use sulphate to oxidise carbohydrates or hydrogen to obtain energy for their metabolism, so there are circumstances in which ΔG is -ve for a reaction scheme in which it behaves as an oxidiser, being reduced to H2S in the process.  But as an oxidiser for what we would regard as a chemical fuel, not really. Concentrated (i.e. not ionised) sulphuric acid is another story however: that can be a powerful oxidising agent. 

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You can use sulphates as oxidisers, if the fuel is a strong enough reductant.
The reaction of calcium sulphate and aluminium  is well documented.
I should probably say something like "don't try this at home" but

 

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8 hours ago, exchemist said:

There are micro-organisms that use sulphate to oxidise carbohydrates or hydrogen to obtain energy for their metabolism, so there are circumstances in which ΔG is -ve for a reaction scheme in which it behaves as an oxidiser, being reduced to H2S in the process.

"Eaters of rock"...

Ever heard/read about the polymetallic nodules on the ocean/sea floor? I don't think they precipitated or accreted over time randomly.

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2 hours ago, NTuft said:

"Eaters of rock"...

Ever heard/read about the polymetallic nodules on the ocean/sea floor? I don't think they precipitated or accreted over time randomly.

Enlighten me. 

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5 hours ago, John Cuthber said:

You can use sulphates as oxidisers, if the fuel is a strong enough reductant.

That goes contrary to my understanding of oxidizers and reducers, John.
Will they spontaneously react, as Sodium and Chlorine, Nitric acid and Hydrazine, or even Lead and Sulphuric acid ?
Or must an initial energy be added, as in your posted video, to get past the potential barrier ?

Please explain the reaction process.
( my last Chem course was Gr.13 in 1976-77 )

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1 hour ago, exchemist said:

Enlighten me. 

We're going off-topic, but, it looks like a job by:

https://handwiki.org/wiki/Biology:Dissimilatory metal-reducing microorganisms

Quote

Dissimilatory metal-reducing microorganisms are a group of microorganisms (both bacteria and archaea) that can perform anaerobic respiration utilizing a metal as terminal electron acceptor rather than molecular oxygen (O2), which is the terminal electron acceptor reduced to water (H2O) in aerobic respiration.[1] The most common metals used for this end are iron [Fe(III)] and manganese [Mn(IV)], which are reduced to Fe(II) and Mn(II) respectively, and most microorganisms that reduce Fe(III) can reduce Mn(IV) as well.[2][3][4] But other metals and metalloids are also used as terminal electron acceptors, such as vanadium [V(V)], chromium [Cr(VI)], molybdenum [Mo(VI)], cobalt [Co(III)], palladium [Pd(II)], gold [Au(III)], and mercury [Hg(II)].[1]

something weird about their cytochrome c's... CymA and TorC genes or something.

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Quote

 

  13 hours ago, John Cuthber said:

You can use sulphates as oxidisers, if the fuel is a strong enough reductant.

That goes contrary to my understanding of oxidizers and reducers, John.
Will they spontaneously react, as Sodium and Chlorine, Nitric acid and Hydrazine, or even Lead and Sulphuric acid ?
Or must an initial energy be added, as in your posted video, to get past the potential barrier ?

Please explain the reaction process.
( my last Chem course was Gr.13 in 1976-77 )

 

Are we talking about anhydrous sulfates? What ones? What is a strong enough reductant? Don't say aluminium.

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10 hours ago, NTuft said:

We're going off-topic, but, it looks like a job by:

https://handwiki.org/wiki/Biology:Dissimilatory metal-reducing microorganisms

something weird about their cytochrome c's... CymA and TorC genes or something.

OK I see what you mean. These nodules contain transition metal ions in low oxidation states, apparently, e.g. Mn (II).  

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Posted (edited)
20 hours ago, chenbeier said:

Why not?

Alkali or earth alkali sulfates can only reduced with aluminium, calcium, sodium and others.

Al uminum! 🤣

j/k. we'll defer to Humphry Davy 's choice about that.

Edited by NTuft
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Posted (edited)
5 hours ago, NTuft said:

Al uminum! 🤣

j/k. we'll defer to Humphry Davy 's choice about that.

Davy, and others at the time, seem to have used a variety of names, Davy initially proposing alumium. Subsequently both aluminium and aluminum were used, both by him and others, though it's true did write a textbook using the aluminum spelling: https://en.wikipedia.org/wiki/Aluminium#Etymology

It seems to have been a choice by a North American called Noah Webster, when compiling his eponymous dictionary in 1828, that settled the American spelling.

Edited by exchemist
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On 5/9/2022 at 6:48 AM, NTuft said:

What is a strong enough reductant?

Cow manure.

Seriously- The farmers used crushed old plasterboard as "bedding" for cattle. (It was cheap)

That led to a dangerous enhanced production of hydrogen sulphide in slurry pits.

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Posted (edited)
On 5/8/2022 at 11:13 PM, chenbeier said:

Why not?

Alkali or earth alkali sulfates can only reduced with aluminium, calcium, sodium and others.

On 5/9/2022 at 1:23 AM, exchemist said:

e.g. Mn (II).

 

I can't always keep it straight myself, so:

Oxidizing agent:

https://handwiki.org/wiki/Chemistry:Oxidizing agent

Quote

An oxidizing agent (also known as an oxidant, oxidizer, electron recipient, or electron acceptor) is a substance in a redox chemical reaction that gains or "accepts"/"receives" an electron from an reducing agent (called the reductant, reducer, or electron donor). In other words, an oxidizer is any substance that oxidizes another substance. The oxidation state, which describes the degree of loss of electrons, of the oxidizer decreases while that of the [Ed.:error here: increases] [reductant] increases; this is expressed by saying that oxidizers "undergo reduction" and "are reduced" while reducers "undergo oxidation" and "are oxidized". Common oxidizing agents are oxygen, hydrogen peroxide and the halogens.

In one sense, an oxidizing agent is a chemical species that undergoes a chemical reaction in which it gains one or more electrons. In that sense, it is one component in an oxidation–reduction (redox) reaction. In the second sense, an oxidizing agent is a chemical species that transfers electronegative atoms, usually oxygen, to a substrate. Combustion, many explosives, and organic redox reactions involve atom-transfer reactions.

 

https://handwiki.org/wiki/Chemistry:Reducing agent

Reducing agent:

Quote

A reducing agent (also known as a reductant, reducer, or electron donor) is an element or compound in a redox chemical reaction that loses or "donates" an electron to an electron recipient (called the oxidizing agent, oxidant, oxidizer, or electron acceptor). In other words, a reducer is any substance that reduces another substance. The oxidation state, which describes the degree of loss of electrons, of the reducer increases while that of the oxidizer decreases; this is expressed by saying that reducers "undergo oxidation" and "are oxidized" while oxidizers "undergo reduction" and "are reduced". Thus, reducing agents "reduce" oxidizers by reducing (decreasing) their oxidation state while oxidizing agents "oxidize" reducers by increasing their oxidation state.

To clarify, in a redox reaction, the agent whose oxidation state increases, that "loses/donates electrons", that "is oxidized", and that "reduces" is called the reducer or reducing agent, while the agent whose oxidation state decreases, that "gains/accepts/receives electrons", that "is reduced", and that "oxidizes" is called the oxidizer or oxidizing agent. A reducing agent is thus oxidized by an oxidizer when it loses electrons that are gained by this oxidizing agent; this oxidizing agent is itself simultaneously reduced by the reducer.

In their pre-reaction states, reducers have extra electrons (that is, they are by themselves reduced) and oxidizers lack electrons (that is, they are by themselves oxidized). A reducing agent typically is in one of its lower possible oxidation states and is known as the electron donor. Examples of substances that are commonly reducing agents include the Earth metals, formic acid, oxalic acid, and sulfite compounds. For example, consider the overall reaction for aerobic cellular respiration:

C6H12O6(s) + 6O2(g) → 6CO2(g) + 6H2O(l)

The oxygen (O2) is being reduced, so it is the oxidizing agent. The glucose (C6H12O6) is being oxidized, so it is the reducing agent.

In organic chemistry, reduction usually refers to the addition of hydrogen to a molecule, though the aforementioned definition still applies. For example, the oxidizing agent benzene is reduced to cyclohexane in the presence of a platinum catalyst:

C6H6 + 3 H2 → C6H12

Historically, reduction referred to the removal of oxygen from a compound, hence the name 'reduction'. An important example of this phenomenon occurred during the Great Oxidation Event, in which biologically−produced molecular oxygen (dioxygen (O2), an oxidizer and electron recipient) was added to the early Earth's atmosphere, which was originally a weakly reducing atmosphere containing reducing gases like methane (CH4) and carbon monoxide (CO) (along with other electron donors) and practically no oxygen because any that was produced would react with these or other reducers (particularly with iron dissolved in sea water), resulting in their removal. By using water as a reducing agent, aquatic photosynthesizing cyanobacteria produced this molecular oxygen as a waste product.[2] This O2 initially oxidized the ocean's dissolved ferrous iron (Fe(II) − meaning iron in its +2 oxidation state) to form insoluble ferric iron oxides such as Iron(III) oxide (Fe(II) lost an electron to the oxidizer and became Fe(III) − meaning iron in its +3 oxidation state) that precipitated down to the ocean floor to form banded iron formations, thereby removing the oxygen (and the iron) [Ed.: stories! lies! alcoa! They hid a bunch of Davy and Faraday's stuff, I bet. They were galvanizing boat hulls w electrolysis ffs. bwahaha...]. The rate of production of oxygen eventually exceeded the availability of reducing materials that removed oxygen, which ultimately led Earth to gain a strongly oxidizing atmosphere containing abundant oxygen (like the modern atmosphere).[3] The modern sense of donating electrons is a generalization of this idea, acknowledging that other components can play a similar chemical role to oxygen.

 

Edited by NTuft
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Posted (edited)
44 minutes ago, John Cuthber said:

Cow manure.

Seriously- The farmers used crushed old plasterboard as "bedding" for cattle. (It was cheap)

That led to a dangerous enhanced production of hydrogen sulphide in slurry pits.

I imagine that would be due to the presence of sulphate-reducing micro-organisms in the manure, wouldn't it?

14 minutes ago, NTuft said:

 

I can't always keep it straight myself, so:

Oxidizing agent:

https://handwiki.org/wiki/Chemistry:Oxidizing agent

 

https://handwiki.org/wiki/Chemistry:Reducing agent

Reducing agent:

 

I find the concept of oxidation states helps. The higher the +ve number, the more oxidised the species, which signifies the removal (whether real or notional) of more electrons.  

Edited by exchemist
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23 minutes ago, exchemist said:

I find the concept of oxidation states helps. The higher the +ve number, the more oxidised the species, which signifies the removal (whether real or notional) of more electrons.

+ve number?

I think I understand the same. Higher positive number meaning net charge higher due to fewer electrons (-) in the balance.

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6 hours ago, NTuft said:

+ve number?

I think I understand the same. Higher positive number meaning net charge higher due to fewer electrons (-) in the balance.

Sort of, yes. So it gives you the idea of losing electrons during oxidation. The caveat (sorry if this is teaching grandmother to suck eggs) is oxidation state is not the same as the charge on an actual ion. It's a book-keeping convention that ascribes a number and sign according to what would apply if a given compound were fully ionic. So for example in SiO2, which of course is covalent rather than ionic, Si has an oxidation state of +4 and O has one of -2.  

And in the present case of the sulphate ion, S has an oxidation state of +6 and the 4 O atoms -2 each. Sulphate reducing bacteria can take S from +6 all the way down to -2 in H2S. 

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