ChemE.info – Chemical Engineering Consulting

A lot of process design is cut and dried but the fun part is finding creative ways to change or improve the process.

I worked for a couple of years on the CUPROSUL process. This process used copper sulfate to scrub H2S out of various gas streams (the principal application was scrubbing geothermal steam). In the original process concept, the sulfur contain in the H2S ended up as ammonium sulfate. It was originally hoped that the ammonium sulfate could be sold as fertilizer. Unfortunately, ammonium sulfate is not widely used as fertilizer in the developed world…

And even worse, the ammonium sulfate that resulted from using the process to scrub geothermal steam contained contaminants that were unacceptable in a fertilizer. This meant that there was no market for the ammonium sulfate and that one would probably have to pay for its disposal.

So, what to do? I had done a literature survey on the reactions used to regenerate the copper sulfate from the copper sulfide produced in the scrubber. From reading the articles it was evident that scrubbed sulfur was briefly present as elemental sulfur in the stirred tank regen reactors but that it was quickly oxidized to the sulfate given the rather severe temperature and oxygen levels.

It occurred to me that, if we could somehow protect the elemental sulfur from further oxidation, we might be able keep it in its elemental form rather than end up with the sulfate form. So I went looking for good sulfur solvents that were immiscible in water and were poor solvents for oxygen. It turns out there are quite a number (including olive oil) but I decided to run an experiment with a chlorinated hydrocarbon which we had in the lab and that seemed to have the desired solvent characteristics.

I took a quantity of the aqueous copper sulfide slurry produced by the scrubber and I oxidized it in an agitated beaker in the presence of the chlorinated solvent. The slurry eventually disappeared and I stopped the agitation; allowing the two solvents to separate. I then decanted the chlorinated solvent into a pan and left it to evaporate in a fume hood. The next morning the pan was dry and covered in sulfur crystals.

That was very satisfying but, of course, we still had to look at the economics of an elemental sulfur by-product. A first pass analysis showed that elemental sulfur by-product did look more promising than sulfate but that the impact would vary by region. The cost of sulfur varies quite widely around the world. Some areas have vast amounts of mineral sulfur that is cheap to mine. Other areas have widely used processes that produce elemental sulfur as a by-product.

So the process modification was not a complete homerun but it did offer the prospect of changing the by-product produced to suit the local market.

There is a tendency to move process development to the next scale prematurely. Since the cost increases substantially with each increase in scale, it makes much more sense to study the process as much as possible at the lower scale.

I have seen organizations jump the gun several times in my career. For one thing, it is never a good idea to move to the next scale if you do not understand the results you are getting at the current scale. I have seen an organization increase from pilot scale to commercial scale even though they were finding that their pilot scale experiments were not repeatable. Needless to say, they then demonstrated their lack of process understanding at the larger scale… At vastly greater expense.

Another example occurred back in the late 70’s when I worked in Polaroid’s department T-22. My role was to scale up processes that the photo and organometallic chemists developed in the lab:

One of chemists asked me to do a pilot-scale production run of a new “receiver layer” system that he had been working on. I asked him to describe how he created the system in the lab and he listed various steps, one of which was the addition of a ferric salt followed by what he called an “incubation period.”

That raised my eyebrows a bit since I didn’t recall that bit of jargon from my reaction kinetics classes and I asked him what he meant by “incubation period.” He was a bit sheepish about it but said that he was adding the ferric salt as a mild oxidizing agent and that it took several minutes before the oxidation that he was looking for was complete.

Since he was using an open beaker, I asked him if he was sure that the ferric was doing the oxidation and not oxygen from the ambient air. I suggested that the “incubation period” might simply be the time required for oxygen to transfer from the room air into the agitated beaker. He was a bit taken aback by the idea but felt that we should test it.

So we reran his laboratory experiment in a narrow mouth flask and sparged the mixture with nitrogen to exclude the room air. We found that we could leave the system indefinitely and that his oxidation would never occur. This confirmed that the ferric had no effect and that the ambient oxygen was the key to his reaction.

The point I am making is that we confirmed this by doing additional lab-scale work with quite small costs in terms of equipment, chemicals, and manpower. We would inevitably have figured out the ferric vs O2 issue at the pilot scale but it would have cost us thousands of dollars more and at least a couple of weeks of time. Once we realized the role of gaseous oxygen, I built a controlled-environment lab setup with a gas manifold that allowed us to meter in the oxygen from a gas cylinder. This allowed us to control the oxidation very precisely and was the approach we used when we finally did move to pilot-scale production.