Crystallization Rate: Prep Room Friday

Before the Christmas break, we had two full classes of Y11 students making clean, dry copper sulfate. A common practical (and a required one on most, if not all, GCSE exams), students LOVED making these beautiful blue solutions – but oh my is there lots of waste.

An evaporating basin, with edges coated with green copper sulfate crystals. There is a small pool of concentrated copper sulfate solution at the bottom.

Rather than waste the resultant solutions, I decided to recrystallise them – especially given that half of the solutions were a gross deep green, after being left to crystallise on the windowsill for a week.

For the uninitiated, recrystallization is the process of dissolving an impure solid and then extracting pure solids from the super-concentrated solution. The solvent you choose must be one in which the solid is only partially soluble, so if you cool the solution, the solid with crystallize out. In theory, these crystals are pure and can be rinsed to give clean, pure crystals.

Their dark green solutions were dissolved into two conical flasks of water, and a few drops of sulfuric acid were added to both, to ensure they stayed as sulfate complexes. Then begins the heating – trying to evaporate off enough water to concentrate the solutions, meaning they’ll crystallise fast, on ice.

I could not, for the life of me, figure out why the student solutions went green. It wasn’t because the students used tap water, forming copper chloride, the only real culprit I could dream of. The addition of more water seemed to bring the cyan back, so my guess is partial dehydration of the complex?

Even after I dissolved student samples, the solutions started turning green – and again, not evenly! So many of the initial student samples were green, that there’s no correlation between the before and after.

This evaporation step took an age – at least two full afternoons. But in the end, they were spread between a few different cooling methods.

Two conical flasks containing copper sulfate solution and a thermometer, blue-green on the left and blue on the right.

Rate of crystallisation is dependent on many things, but the major one here is temperature. As the temperature of the solution drops, the solubility of the copper sulfate in water also drops, and the solid starts to crystallise out. The longer and slower you let the solution cool, the longer you give the crystals to form, and (in theory) the larger the crystals you get.

As such, I tried a few methods: evaporating basins in an ice bath (the usual method) and scratched Petri dishes on an ice block (the new one). Plastic Petri dishes are far thinner, have a wider surface area and can be scratched, creating a nucleation point (or starting point) for the crystals and a far faster drop in solution temperature. Together, these should form crystals fast enough for students to see quickly, compared to the overnight for an evaporation basin.

And hey presto – we saw crystals within near-instantly! It took far longer for evaporation basins, which sat for aroun half an hour. Petri dishes gave fast, powdery, small crystals, compared to the slower but larger crystals of an evaporating dish.

The left-over filtrate from these crystallised solutions was again concentrated by evaporation, and this time allowed to evaporate over the two-week break in an evaporation basin – and the crystals formed were the best of them all.

Viewed under the microscope, they are beautiful. See below – with petri dish crystals top left, iced evaporating basin crystals bottom left, and slow evaporating basin crystals on the right.

In short, the longer the crystallisation process, the larger (and smoother) the crystals. While this isn’t included on many specifications nowadays, this is the same concept behind rock formation: the longer the rock has to cool from magma, the larger the crystallites within the rock, giving air-cooled rocks far larger crystals than ocean-cooled ones.

Small, irregular blue crystals of

Feel free to show these pictures to students! Then perhaps they can guess how the chemical supply company purifies their copper sulfate, again under the microscope (on the left)?

Now, these samples are ready to be redissolved and used – perhaps as an electrolyte for electrolysis, or preparation of fresh Biuret solution.

As a Chemistry specialist technician with experience in teaching, I’d love to work on more accessible practicals with purpose. If you have ideas, let me know here.