Introduction
So at some point in everyone's chemistry education they will end up doing an experiment that produces hydrogen gas and they end up filling a balloon with hydrogen and igniting it... Or so I thought. See this is a cheap yet entertaining and relatively safe experiment that any high school science or chemistry teacher could conduct but apparently most of my friends that I have talked to (even some going to the same high school but getting different chemistry teachers) have never seen this reaction occur.
Note for teachers, this creates a small fireball and will make a thud noise about as loud as a balloon popping. My chemistry teacher did this in the classroom without issues but it can also be done outside.
So what I did was I started making some hydrogen gas and I filled some balloons with it and I took a lighter to the balloons (barbecue lighter but I have done it with just a handheld BIC lighter). I also recorded it and have 3 recordings combined into a single video embedded below. However, right now it is time for some chemistry.
Chemistry
Warning this section contains the use of math and equations, feel free to skip to the next section if that doesn't interest you.
Reaction: 2H2(g)+O2(g)→2H2O(g)+energy
This reaction looks pretty straightforward, I mean besides the energy term which I will explain later. What it says is that when hydrogen gas and oxygen gas combine they create water vapour. This reaction also releases a energy to the surrounding environment as heat, light, and sound and I am going to go through the enthalpy calculation to determine roughly how much heat.
Now to calculate the energy released from a reaction you have a few options... First you can go the calorimetry route using the equation ∆H=mc∆T which is all well and good but it isn't very good for theoretical applications. Second is that you can use the bond energies to determine the energy released from the reaction. The third way is to use the heat of formation to determine the energy released from the reaction. I am going to estimate the energy released from both heat of formation and bond energy
Heat of Formation
| Chemical | Heat of Formation (kJ/mol) |
|---|---|
| H2(g) | 0 |
| O2(g) | 0 |
| H2O(g) | -241.820 |
Now the above equation states that the energy released in a reaction is equal to the sum of the heat of formation of the products of the reaction subtract the sum of the heat of formation of the reactants. In practice this would be calculated as follows:
∆Hp = 2(mol H2O) * -241.820 = -483.640
∆Hr = 2(mol H2) * 0 + 1(mol O2) * 0 = 0
∆Hf = ∆Hp - ∆Hr = -483.64-0
∆Hf = -483.64kJ
So this means if we have 20 mols of hydrogen that get reacted with oxygen (assuming 100% reaction rate) then we will get 4836.4kJ of energy released to the surrounding area (negative energies here mean that energy is released and that the reaction is exothermic. Positive energies mean the reaction is absorbing energy and is endothermic).
Bond Energy
| Chemical | Bond Energy (kJ/mol) |
|---|---|
| H2 | 436 |
| O2 | 498 |
| H2O | 926 |
For the table, the energy of a hydrogen-hydrogen bond is 436kJ/mol, and a hydrogen-oxygen bond is 463kJ/mol and oxygen-oxygen bond is 498kJ/mol. However water (H2O) contains 2 hydrogen-oxygen bonds meaning that it gets doubled. Also when it comes to calculating the total energy released I will be multiplying by the number of mols of each chemical meaning we will effectively double both the water and hydrogen bond energies for the final calculation.
∆Hr = 2(mol H2) * 436 + 1(mol O2) * 498
∆Hr = 1370
∆Hp = 2(mol H2O) * 926
∆Hp = 1852
∆Hrxn = 1370 - 1852
∆Hrxn = -482kJ
Discussion
So with method 1 we predicted an energy released of 483.64 kJ and with method 2 we predicted a value of 482 kJ being released. The cause for this discrepancy is simply to do with the tables I was using (especially for the bond energies) not being super precise. Had I used more accurate tables that were more specific then the numbers would be the exact same but ultimately that doesn't really matter here. We now know the theoretical amount of energy released per mol of oxygen consumed in the combustion of hydrogen gas.
Another thing to note is that we could have skipped all these calculations and looked up the enthalpy of combustion of hydrogen and then multiply that by the number of mols of hydrogen consumed. Now if you do look that up you may find sources saying the enthalpy of combustion of hydrogen is around -285 but this is the enthalpy of combustion when liquid water is produced and my calculations are assuming it produces water vapour. If you assume liquid water is produced then you get around -570kJ/mol which is a little more than the 482kJ/mol. The only reason why I didn't mention it above is that enthalpy of formation and enthalpy of combustion are really the same thing but because enthalpy of formation has more use cases I went with that.
Producing Hydrogen Gas
Electrolysis of Distilled Water
If you look to the left you will see a table of half reactions and standard potential. You do not really need to know how to read the table, what does matter is the half reaction of O2+4H++4e-→2H2O. This reaction has a standard potential of 1.23V meaning to do the reverse reaction where we produce hydrogen gas we need to supply at least 1.23V to it to break water into hydrogen and oxygen. Now using some basic electrical knowledge with Ohms law we have V=IR where I is the current, V is the voltage, and R is the resistance... The point of this is that when we increase the current the faster the reaction will take place and since current is directly proportional to voltage and inversely proportional to resistance we can increase current by either increasing the voltage or reducing the resistance.I bring this up because a lot of people will think "Hey we could just dissolve some ions in the water and it will reduce the resistance and salt is a commonly available ion." but I am making this warning here because adding table salt to the mixture can do something anomalous where you will start producing chlorine gas. As such this is commonly called the chloride anomaly. I have linked a couple of sources that go more in depth and eventually I will make a post going more in depth on the chloride anomaly. However when you read the standard half reaction table it would appear like the hydrogen and oxygen production would be preferred over the chlorine production (due to chlorine production having a higher standard potential) and this is true... But this table is leaving out a lot of information. I just wanted to bring this up before someone makes a concentrated brine solution for hydrogen synthesis and then wondering why chlorine gas is produced instead.
Zinc metal and Hydrochloric Acid
To predict this reaction you need to know a couple things . First zinc has an oxidation state of +2 meaning you can bind 2 chlorine (-1) atoms to it in order to make it stable. Second is the electronegativity of the chemicals...
| Chemical | Electronegativity |
|---|---|
| Zn | 1.6 |
| H | 2.2 |
| Cl | 3.2 |
Now things tend to like to bind to other molecules that have as large of a difference in electronegativity. So the difference between hydrogen and chlorine is 3.2-2.2=1.0 where the difference between zinc and chlorine is 3.2-1.6=1.6... 1.6>1.0 therefore the chlorine will prefer to react with the zinc. This is generally speaking of course, like its more complicated when you get into higher level chemistry but for simple reactions this can be used to predict a number of simple reactions.
Aluminium and Sodium Hydroxide
Alright I am just going to say something. Researching this reaction is weird because a lot of people will argue that nothing will happen. Now, obviously it does in fact happen where you add solid aluminium to an aqueous sodium hydroxide solution and you get Sodium Aluminate precipitating out and hydrogen gas forming. So here it the reaction that occurs:
It is important to note that without water present then aluminium and sodium hydroxide do not react but some people assume that this will remain consistent with water present not realizing that water changes what occurs. Further I do not feel confident in my chemistry knowledge or ability to tell you exactly why this is the chemical that ends up forming in this reaction especially since there is a very similar reaction shown below where water is still present but because it isn't in solution it changes how the reaction occurs. The product produced is also called sodium aluminate.
Again, why these reactions occur as they do, I do not feel comfortable in my knowledge to try to explain. However if someone with more chemistry knowledge wants to take a crack at it as a reply, be my guest. Here is a wikipedia article on sodium aluminate for further reading.
Other
Quite honestly I could go on for a really long time with different methods of producing hydrogen. Honestly too, I could probably write an entire post about each of the methods already covered (including the sodium aluminate one, I am trying to keep this post somewhat short). Instead I am just going to drop some links here cause lazy.
potassium hydroxide and aluminium
wikipedia list
Video
I produced the hydrogen by reaction sodium hydroxide and aluminium from unrecycled pop cans. Only bring this up because I have no doubt, as is every time I have read about this reaction, some hyped chemistry student will say "no reaction will occur" after they spend 30 seconds looking at a redox table. I assure you, the reaction occurs.
Conclusion
Now I just want to reiterate why I think teachers especially should look into doing a hydrogen demonstration. First you can use it as a method to teach a number of different concepts from electrolysis, electrochemistry, combustion, enthalpy, calorimetry, molecular bond energies, basic stoichiometry, electronegativity, and more. Second, not only does it act as a method to educate these topics but its a fun demonstration with relatively minimal risk (the explosion isn't super loud as its a deflagration, aka subsonic... And it is easy to control to make sure you don't cause injury). Another benefit to this demonstration is that it is also extremely cheap and easy to pull off to the point that most people can do it with basic household chemicals just lying around.
Now if you aren't an educator then maybe you are a parent or know someone with kids or you just want to impress your friends... Now you can grab a balloon and go all Walter White and mix some chemicals together, fill the balloon with hydrogen and set it off and wow them. Honestly I have done stuff like this as party tricks and everyone has a good time.
Sources
Standard Enthalpy of Formation
Table of Bond Energies
Chloride Anomaly Biology Forums
Chloride Anomaly Chemistry Forums
Standard Potential Image Source
Table of Electronegativities
Sodium Hydroxide and Aluminium with an answer saying nothing happens