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megalomania
June 20th, 2002, 02:05 AM
I have found some information about using microreactors to simulate the conditions of the Ostwald process for the production of nitric acid. Microreactors are small microchip sized chemical reactors that have conditions similar to their industrial counterparts. I find this data relevant because it I am also trying to build a microreactor, although bigger.
First we consider the reactions we are going to encounter:

4NH3 (ammonia) + 5O2 (oxygen) --> 4NO (nitric oxide) + 6H2O (water) - 904 KJ
4NH3 (ammonia) + 6NO (nitric oxide) --> 5N2 (nitrogen) + 6H2O (water) - 1816 KJ
2NH3 (ammonia) --> N2 (nitrogen) + 3H2 (hydrogen) + 93 KJ

Experimental tests varying only the initial reaction temperature (pressure, 1 atm held constant) show that equations 1 and 2 are very competitive with each other at 400 C, whereas equation 3 can largely be neglected. For equations 1 and 2 only, the amount of ammonia consumed is 70%, and the conversion into NO is 1.9%. Factoring in equation 3, the amount of ammonia consumed is 72% and the yield is 1.9%.
Increasing the initial temperature up to 1400 C shows that equation 2 becomes negligible, and equation 3 competes with equation 1. At this temperature we see 84% of the ammonia being consumed and 46% being converted into NO.
Varying the temperature over a wider scale, a clearer picture of optimal temperature conditions can be seen. Keep in mind that at low temps, equation 2 dominates, and at extreme temps, equation 3 dominates. We wish to find the best temperature range where equation 1 dominates, and the yield of NO is highest:

Temp: 200 C - 400 C – 600 C - 800 C – 1000 C – 1200 C – 1400 C – 1600 C
Yield: 0.5% -- 2% ----- 26% --- 63% --- 74% ----- 63% ----- 46% ----- 31%

Notice that at 1000 C we get the highest yield of NO. It is the range, between 800 C to 1200 C, that we want to shoot for. Of course the reaction itself will generate heat, and that heat will affect the equilibrium of the reaction. Furthermore, the catalyst chamber will be at different temperatures through its length. The temperature of the reaction will increase until equilibrium is reached, and there will be different reactions temperatures from start to finish. The real world data on this are as follows:

Init temp:- 200 C -- 400 C --- 600 C --- 800 C -- 1000 C - 1200 C - 1400 C - 1600 C
Exit temp: 1406 C - 1566 C - 1637 C - 1695 C - 1744 C - 1790 C - 1838 C - 1894 C
Real Yield: 37% --- 39% ----- 37% ---- 36% ---- 33% ----- 39% ----- 25% ----- 21%

Now this shows us there is little real difference between any reaction temperatures, except at higher levels. Factoring in pressure drop as the reaction proceeds through the catalyst chamber, we get a slightly different picture on our yield. The pressure drop is the result of reactants being consumed thus decreasing the interaction between ammonia and oxygen. This does not mean the real pressure changes. Don’t tell me in a whiny voice that the number of moles of substance could increased thus increasing the volume of gas causing a net increase in pressure, because I can see that. This pressure drop is using a chemical engineers definition.

Init temp:- 200 C -- 400 C --- 600 C --- 800 C -- 1000 C - 1200 C - 1400 C - 1600 C
Exit temp: 1288 C - 1484 C - 1564 C - 1628 C - 1682 C - 1732 C - 1785 C - 1848 C
Real Yield: 26% --- 34% ----- 35% ---- 33% ---- 30% ----- 26% ----- 22% ----- 18%

Now we can see that a rise in temperature is actually bad for NO production. How does all this relate to a means of actually producing nitric acid on a scale that is useful? These results show that the temperature of the system is the most important factor. The remaining data from these experiments shows that the actual pressure drop is about 50% to 0.5 atm, changing the length of time that the ammonia is in contact with the ammonia is best at 0.0003 s, and that the minimum ratio of oxygen to ammonia is 2:1, but these factors change the actual yield very little.
As you can see, there is a big difference between starting this at 200 C versus 600 C. It would seem to me that once this reaction gets going, any external heat should be removed, or reduced, so the reaction can commence. It also occurs to me that some means of measuring the temperature of the system might be beneficial. The final recommendation of these experiments is to maintain a temperature at 1000 C.

There are a few insightful guidelines that can be gleaned from this data. First we will start with the ratio of air to ammonia. You may tell from equation 2 that the presence of too much ammonia may contribute to that reaction. These experiments have demonstrated this is the case. It seems to me that you can’t really go wrong with increasing the flow of oxygen. There is only a slight change in the % yield by changing these ratios. All of the data I have indicates that an optimal ratio of oxygen to ammonia is around 7:1. While this will vary from system to system, it is not really an important factor, at least at this time. Ramping up the ratio of oxygen would be the best way to start if you have no flow control, such as my system. I may include some pressure valves in my future designs to assist in the more precise control of both the volumes and the pressure of the gas.

That leads me into the second factor, the pressure of the system. Again, there is little change in the yields based on the pressure, but if you can increase the pressure by all means do so. I have data that says a range from 1 to 100 atm is equally acceptable. Industrially they have both low and high pressure systems. Following their example and applying it to our needs, we need not concern ourselves with pressure. We could get by with an open system at atmospheric pressure just as well as sealing it up to increase the pressure a few atms.

Third is our residence time. This is tricky to calculate, but easy to determine experimentally. Every system will have different conditions depending on the catalyst, the length of the catalyst chamber, the fluid flow around the catalyst bed, and the speed of the fluid going through it. One could regulate the speed of the gasses going in by pressurizing the system or increasing the gas flow. As the gas speeds up, the residence time decreases, and the amount of nitric acid changes. Experimentally you can speed up the gas until you start to get unreacted ammonia detected as ammonium nitrate. Slow it down too much and you get no acid at all. Somewhere in the middle it will be optimal for you. Again, every catalyst is different, as is every system when you stuff the catalyst in your reaction chamber.
I have just begun to think about how to ‘detect’ the presence of unreacted ammonia. The original patent literature assumes you smell it, but that seems rather dangerous to me. I hate the smell of ammonia, and I dislike breathing nitrogen dioxide gas. The NO2 gas may be heavy enough to sink, while the ammonia is detectable, but this seems so unscientific. One could titrate the nitric acid and weigh it on an analytical balance to see how much extra ammonium nitrate there is, but for how many people would this be practical? Maybe just boiling off all the liquid would work. Anybody have any good ideas?

Our last, and most important, factor is temperature. I see now that the presence of so much water vapor in my earlier experiments was very bad indeed. Unless you can heat this puppy up to 600 C to start with, you will be wasting ammonia and time. The good news is that as the reaction proceeds it heats itself, and so you should be able to reach temperatures of 1000 C even if you start at a low 200-300 C. In fact, since the lower temperatures favor the second equation, and since that equation produces twice as much heat energy, the amount of wasted ammonia may be minimized. The one set of data I do not have is how long it takes the system to reach its equilibrium temperature.
I have a pyrometer laying around somewhere; I hope it still works. I actually had some data that showed the equilibrium temperature of this system, but I lost it. I think it is 600-800 C, which is pretty good. Experimentation with temperature monitoring would seem in order here.

rikkitikkitavi
June 20th, 2002, 08:00 AM
One idea about detecting NH3:

Any NO2 formed decomposes into O2 and NO above 300 C.

This means that the hot gases after the catalysts contain little NO2.

NO has very low solublity in water.

NH3 is very soluble in water. The dissolving process of NH3 is faster than NO + O2 = > NO2 and the NO2-dissolving process.

My idea:

letting the hot reaction gases after the catalyst enter a scrubber (washing column) any NH3 resident is dissolved , and the short resident time means that little NOx is lost this way, and the sideeffect is that the gases leaving the scrubber are cooled, which is good for our synthesis of NO2.
Most of the heat generated in the catalyst is cooled away in the scrubber. But water has very high heat capacity, which is important since NH3 solubilty goes down when temperature goes up.

Measuring the pH(f e x with a pH-meter) at the output could give a indication of how much free NH3 is present.

/rickard

vulture
July 8th, 2002, 05:45 AM
The thing I'm worried about is reaction 3 competing with reaction 1, hot NOx, O2 and H2 in a reactor would very likely cause an explosion.

<small>[ July 08, 2002, 04:45 AM: Message edited by: vulture ]</small>

megalomania
July 9th, 2002, 01:58 AM
I did have some data on that, and there is that possibility. However, it cam be minimized because the temperatures required to do this are much higher than the system is run at, and there is a rather narrow % gas mix to get explosion levels. By diluting the gas with enough air, you can bring down the risk. Proper temperature control will also e,iminate the risk. If allowed to proceed with enough insulation, a system can just about double its heat to temps able to detonate the gas mix. This problem is only a problem on a very large scale where thermal dissipation is more difficult. On the bench scale (and in microreactors) you can get it that hot because you could not pour enough reactants into the system of this size.

If you allow the gasses to accumulate in a confined space for a long enough time, sure you can get an explosion, but this is very unlikely. In order for this to happen you would need to just let it vent into the open in a tiny room. Why you would want this much heat in a tiny room with it just spewing nitrogen dioxide gas all over is beyond me. If this is the case you are obviously not watching your experiment and you have other problems like nitrogen dioxide spewing all over.

THe_rEaL_dEaL
February 19th, 2003, 10:53 AM
A question for the multiple flask process to ensure maximum NO<sub>2</sub>/NO conversion to Nitric.

Is pollyproplyene plastic resistant to concentrated nitric acid?
If so I have heard that the flexible, black reticulation piping sold in hardware and retic stores is polly pro. Also soft drink bottle plastic used in the 600ml variety is supposedly polly pro.

This could be very useful plastic. :)

<small>[ February 19, 2003, 09:58 AM: Message edited by: THe_rEaL_dEaL ]</small>

Anthony
February 19th, 2003, 02:37 PM
Plastic labware is made of polypropylene, and IME is HNO3 resistant.