I do not understand oxidizers and what they do

  • #1
So I was doing some research on the Beirut explosion and the consensus is it was caused by around 3000 tons of ammonium nitrate. I always knew things like ammonium and potassium nitrate were explosive but I never knew why. So apparently they are called oxidizers. First off im confused because I thought oxidation had to do with metals rusting not explosions. Second I read that the oxidizer isn't the thing that blows up or is burned, its the thing that makes the thing being burned burn faster. So what is the truth, did the Ammonium nitrate blow up or did it facilitate the explosion of something else in that Beirut warehouse? Also, I looked up what an oxidizer is and its something that can take electrons from something else. Isn't that what an acid is, a proton donor or by some definitions (lewis dot I think) an electron acceptor, I need that cleared up, because then what makes an acid and an oxidizer different. Now finally, to attempt to understand the underpinnings of the chemistry here, why would the accepting of an electrons (what an oxidizer does) cause a massive explosion, is this combustion? I don't understand how oxidizers would lead to massive explosions.
 
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  • #2
DaveC426913
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Rust and explosion are merely opposite ends of the scale called oxidation. One is very, very slow; the other is much faster.

There is plenty of explanatory literature out there, especially in the wake of the Beirut explosion. Have you tried Googling it?
 
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  • #3
DaveC426913
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Like many explosives, such as nitroglycerin, all they really do is decompose into gases, such as CO2, water vapour, oxygen and nitrogen. Oh. And heat.

These are gases, and as such, they prefer to expand. They prefer it a LOT. And they are not known for their patience in doing so.
 
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  • #4
Interesting, I still don't understand how the oxidizer in the explosion caused it. What did it oxidize? It seems like the ammonium nitrate just exploded, but that's not the case is it? The ammonium nitrate oxidized something else and stripped its electrons of that thing and then that somehow turned into a gas? So its a catalyst. Im still confused .
 
  • #5
So I did some additional reading, these oxidizers not only take an electron but they always give up an oxygen. Is that correct? Which makes sense in terms of an explosion. Is it the case that the electron has nothing to do with it and the ammonium nitrate acts as a fuel for instant oxygen? correct me if im wrong.
 
  • #7
Vanadium 50
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Rust and explosion are merely opposite ends of the scale called oxidation. One is very, very slow; the other is much faster.

I owned a 1976 Chevrolet Malibu, and can attest that they take about the same time. :wink:

@Nick tringali , it's probably best to think of the destruction as due to a fast chemical reaction that released a lot of energy. Exactly which reaction is less important.
 
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  • #8
chemisttree
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Smoldering ammonium nitrate can decompose into gaseous ammonia (fuel) and various gaseous oxides of nitrogen like nitrous oxide (oxidant). Like a fuel-air bomb on steroids. Another oxide of nitrogen is NO2 <->N2O4 otherwise known as rocket fuel oxidizer. All those oxidizers will combine with the ammonia and generate gaseous products.

A better explanation of the chemistry of the Beirut explosion, IMO.
 
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  • #9
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So I was doing some research on the Beirut explosion and the consensus is it was caused by around 3000 tons of ammonium nitrate. I always knew things like ammonium and potassium nitrate were explosive but I never knew why. So apparently they are called oxidizers.
Potassium nitrate is not explosive - only component of explosives. Ammonium nitrate is a component of explosive but also itself explosive.
First off im confused because I thought oxidation had to do with metals rusting not explosions.
Both are oxidation reactions, as stated before.
Second I read that the oxidizer isn't the thing that blows up or is burned, its the thing that makes the thing being burned burn faster. So what is the truth, did the Ammonium nitrate blow up or did it facilitate the explosion of something else in that Beirut warehouse?
In Beirut, it blew up itself.
Also, I looked up what an oxidizer is and its something that can take electrons from something else. Isn't that what an acid is, a proton donor or by some definitions (lewis dot I think) an electron acceptor, I need that cleared up, because then what makes an acid and an oxidizer different.
Different kinds of products they are trying to form. Like the contrast between chlorine (strong oxidizer but not acid) and hydrogen chloride (strong acid but a weak oxidizer).
Now finally, to attempt to understand the underpinnings of the chemistry here, why would the accepting of an electrons (what an oxidizer does) cause a massive explosion, is this combustion? I don't understand how oxidizers would lead to massive explosions.
Yes. Not all redox reactions are massively energetic, but most of the most energetic reactions are redox reactions. Which is why combustion and explosions are usually redox reactions.
 
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  • #10
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So I did some additional reading, these oxidizers not only take an electron but they always give up an oxygen.

That's the usual case (and where the term "oxidation" comes from). But oxidizers don't need to release oxygen. Here is a simple counter example:

##2Na + Cl_2 \to 2NaCl##

This is a redox reaction without any oxygen. The oxidizer chlorine is reduced from Cl2 to Cl- (it accepts an electron) and the reducting agent sodium is oxidised from Na to Na+ (it releases an electron).

Now you might ask how to see if a reaction is a redox reaction or not. There is a property called "oxidation state". You can assign an oxidation state to every atom of a compound by applying the corresponding rule set. It is increased by oxidation and decreased by reduction. In a redox reaction you have both, increasing and decreasing oxidation states. Compounds, groups or single atoms (depending how deep you want to brake it down) with decreasing oxidation states are the oxidizers and reduction agents are characterized by increasing oxidation states.

In the example above it is clear. The oxidation state of Na increases from ±0 to +1 and the oxidation state of Cl decreases from ±0 to -1. But it can be complicate. Let's take two possible decomposition reactions of ammonium nitrate with the oxidation states above the atoms:

(1) ##\mathop N\limits^{ - 3} \mathop {H_4 }\limits^{ + 1} \mathop N\limits^{ + 5} \mathop {O_3 }\limits^{ - 2} \to \mathop N\limits^{ - 3} \mathop {H_3 }\limits^{ + 1} + \mathop H\limits^{ + 1} \mathop N\limits^{ + 5} \mathop {O_3 }\limits^{ - 2}##

(2) ##\mathop {2N}\limits^{ - 3} \mathop {H_4 }\limits^{ + 1} \mathop N\limits^{ + 5} \mathop {O_3 }\limits^{ - 2} \to 2\mathop {N_2 }\limits^{ \pm 0} + \mathop {O_2 }\limits^{ \pm 0} + 4\mathop {H_2 }\limits^{ + 1} \mathop O\limits^{ - 2}##

The first example is no redox reaction because the oxidation states of all atoms remain unchanged. The second example is a redox reaction because the oxidation states of nitrogen change. But ammonium nitrate is both, oxidizing and reducing agent. You need to brake it down to the atoms to see what is reduced and what is oxidized:

The nitrogen of the ammonium ion NH4+ is oxidized from -3 to ±0,
the nitrogen of the nitrate ion NO3- is reduces from +5 to ±0 and
one oxygen atom is oxidized from -2 to ±0.

That's what most likely happened during the Beirut explosion (with additional equilibrium reactions between nitrogen and oxygen to a mixture of nitrogen oxides).

Is it the case that the electron has nothing to do with it and the ammonium nitrate acts as a fuel for instant oxygen?

The electrons actually have something to do with it. Shifting electrons away from oxygen and one of the nitrogen atoms and toward the other nitrogen atom is what the reaction makes a redox reaction and redox reactions usually (but not always) have a high reaction heat. Ammonium nitrate being both, oxidizer and reducer at once is another important factor because both parts of the reaction are always available at the same place. The third factor has also been mentioned above: The reaction produces a lot of gas that expands due to the reaction heat. That's what the destruction finally comes from.
 
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  • #11
Potassium nitrate is not explosive - only component of explosives. Ammonium nitrate is a component of explosive but also itself explosive.



In Beirut, it blew up itself.
Thank you. So what is the molecular underpinnings of why ammonium nitrate can explode itself but Kno3 cannot?
 
  • #12
That's the usual case (and where the term "oxidation" comes from). But oxidizers don't need to release oxygen. Here is a simple counter example:

##2Na + Cl_2 \to 2NaCl##

This is a redox reaction without any oxygen. The oxidizer chlorine is reduced from Cl2 to Cl- (it accepts an electron) and the reducting agent sodium is oxidised from Na to Na+ (it releases an electron).

Now you might ask how to see if a reaction is a redox reaction or not. There is a property called "oxidation state". You can assign an oxidation state to every atom of a compound by applying the corresponding rule set. It is increased by oxidation and decreased by reduction. In a redox reaction you have both, increasing and decreasing oxidation states. Compounds, groups or single atoms (depending how deep you want to brake it down) with decreasing oxidation states are the oxidizers and reduction agents are characterized by increasing oxidation states.

In the example above it is clear. The oxidation state of Na increases from ±0 to +1 and the oxidation state of Cl decreases from ±0 to -1. But it can be complicate. Let's take two possible decomposition reactions of ammonium nitrate with the oxidation states above the atoms:

(1) ##\mathop N\limits^{ - 3} \mathop {H_4 }\limits^{ + 1} \mathop N\limits^{ + 5} \mathop {O_3 }\limits^{ - 2} \to \mathop N\limits^{ - 3} \mathop {H_3 }\limits^{ + 1} + \mathop H\limits^{ + 1} \mathop N\limits^{ + 5} \mathop {O_3 }\limits^{ - 2}##

(2) ##\mathop {2N}\limits^{ - 3} \mathop {H_4 }\limits^{ + 1} \mathop N\limits^{ + 5} \mathop {O_3 }\limits^{ - 2} \to 2\mathop {N_2 }\limits^{ \pm 0} + \mathop {O_2 }\limits^{ \pm 0} + 4\mathop {H_2 }\limits^{ + 1} \mathop O\limits^{ - 2}##

The first example is no redox reaction because the oxidation states of all atoms remain unchanged. The second example is a redox reaction because the oxidation states of nitrogen change. But ammonium nitrate is both, oxidizing and reducing agent. You need to brake it down to the atoms to see what is reduced and what is oxidized:

The nitrogen of the ammonium ion NH4+ is oxidized from -3 to ±0,
the nitrogen of the nitrate ion NO3- is reduces from +5 to ±0 and
one oxygen atom is oxidized from -2 to ±0.

That's what most likely happened during the Beirut explosion (with additional equilibrium reactions between nitrogen and oxygen to a mixture of nitrogen oxides).



The electrons actually have something to do with it. Shifting electrons away from oxygen and one of the nitrogen atoms and toward the other nitrogen atom is what the reaction makes a redox reaction and redox reactions usually (but not always) have a high reaction heat. Ammonium nitrate being both, oxidizer and reducer at once is another important factor because both parts of the reaction are always available at the same place. The third factor has also been mentioned above: The reaction produces a lot of gas that expands due to the reaction heat. That's what the destruction finally comes from.
Thank you. So you mentioned that an electron was taken from one oxygen and one nitrogen and those 2 were given to another nitrogen. Which nitrogen and oxygen lost the electron in ammonium nitrate and which nitrogen gained it? Furthermore, because ammonium nitrate is comprised of nitrogen in ammonium and a nitrogen in nitrate that gives it the special property to act as a oxidizer and a reducer? What if the warehouse was comprised of KNO3 and not ammonium nitrate, would the situation have turned out differently because KNO3 only has 1 nitrogen can can only act as a component of an explosion and not be the explosion itself unlike ammonium nitrate that can blow up itself?
 
  • #13
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Which nitrogen and oxygen lost the electron in ammonium nitrate and which nitrogen gained it?

Just check the oxidation states. They are reduced by additional electrons and increased by "lost" electrons (the atoms don't turn into ions).

Furthermore, because ammonium nitrate is comprised of nitrogen in ammonium and a nitrogen in nitrate that gives it the special property to act as a oxidizer and a reducer?

It is because there are both, oxidizing and reducing parts in the same compound. There doesn't need to be nitrogen involved. It the case for ammonium nitrate but you can have similar effects with other elements.

What if the warehouse was comprised of KNO3 and not ammonium nitrate would the situation have turned out differently because KNO3 only has 1 nitrogen can can only act as a component of an explosion and not be the explosion itself unlike ammonium nitrate that can blow up itself?

The equation (again with oxidation states) for the thermal decomposition is

##\mathop {\rm K}\limits^{{\rm + 1}} \mathop {\rm N}\limits^{{\rm + 5}} \mathop {{\rm O}_{\rm 3} }\limits^{{\rm - 2}} \to \mathop {\rm K}\limits^{{\rm + 1}} \mathop {\rm N}\limits^{{\rm + 3}} \mathop {{\rm O}_{\rm 2} }\limits^{{\rm - 2}} + {\textstyle{1 \over 2}}\mathop {{\rm O}_{\rm 2} }\limits^{ \pm {\rm 0}}##

Potassium remains at +1. It is just sitting there doing nothing. Nitrogen is reduced from +5 to +3. Oxygen is oxidized from -2 to ±0 and oxygen doesn't want to be oxidized. This reaction is endothermic. That means no explosion with KNO3 (and without reducing agents).
 
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  • #14
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What if the warehouse was comprised of KNO3 and not ammonium nitrate, would the situation have turned out differently because KNO3 only has 1 nitrogen can can only act as a component of an explosion and not be the explosion itself unlike ammonium nitrate that can blow up itself?
Yes. KNO3 is just an oxidant, so it does not blow up alone.
However, a mixture of KNO3 with suitable reducers can and does blow up. Such as coal and brimstone - which with saltpetre make up gunpowder. And often has exploded.
 
  • #15
chemisttree
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Oxidation state of nitrogen in ammonia is -3.
Oxidation state of nitrogen in nitrate is +5.

NH4+NO3- —> N2 + 1/2 O2 + 2H2O

... is the high temperature reaction given in the literature. Can you find anything wrong with this?
 
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  • #16
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In my opinion, oxidation is a concept that is very helpful in some places, but because it is not fundamental, it can be misused. Oxidation has three different meanings: removal of hydrogen, adding oxygen, or losing an electron (as in Fe II -> Fe III). Unfortunately, these meanings are not really equivalent. If you think of methane as a fairly neutral sort of molecule, and you heat it, you progressively get ethene, then acetylene, through removing hydrogen. Technically, oxidation. Continue and you get carbon, one of the stronger reducing agents. The concept works very well for ions, but less so when covalent bonding is present.

If we take ammonium nitrate, the reaction is
NH4NO3 -> 2H2O + N2O
The OH bond is a good bit stronger than the NH bond, and NO3- is not exceptionally stable when removed from the electric field of nearby cations, and while ammonium is a cation, it does not exert a particular strong local electric field. The net result is this transformation is rather exothermic, and the products are gaseous, thanks to the heat. That means a few mL of solid generates over 66 L of gas, thus in a constrained space, generating a large pressure. Unfortunately, ammonium nitrate decomposition is spontaneous under pressure. Apparently it was recognised as explosive in Germany when a large pile of it outside got wet and it was attacked with some sort of pick or jackhammer. If a small sample decomposes and the pressure cannot be released a shock wave goes through the mass and boom.

Basically, the reaction is simply the bonds or electron distribution simply rearranging to go to a lower energy. If you want you can say the hydrogens were oxidised to water, but in this case you could also say water was eliminated and the nitrogen atoms formed a more stable double bond system. In my opinion, when you have independent ions oxidation is a useful concept to explain what is going on, but it is less so in systems and for covalent bonds.
 

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