Analytical chemistry is a strange hobby of mine, and not the easiest one to indulge. Even the most basic equipment for carrying out chemical analyses can be costly, beginning with the requirement of an analytical balance of great reliability and accuracy. (Precision to 0.0001 grams is regarded as necessary, but I'm making do with a cheaper balance of 0.001 gram precision.) But equipment is easy enough to get hold of, if you have the money. What's really getting difficult to find are the chemicals. More and more supply houses are not selling to the general public.

Some chemicals used in analysis need not be very pure. For example, it is never expected that even good grades of sodium hydroxide, used in acid-base titrations, is very certain in its composition. The same is true to a lesser extent of hydrochloric acid, another valuable reagent in volumetric analysis. The trick is to prepare solutions of approximately known concentration and then experimentally determine their exact concentration by using them to titrate carefully weighed samples of substances (almost always solids) of trustworthy composition and purity. This technique is called standardization, and it means that I can use cheap household lye instead of high-grade sodium hydroxide, and hardware store muriatic acid instead of pure hydrochloric acid, as long as I can standardize the solutions prepared from them.

But you still need those pure compounds, the primary standards. A real chemist can easily get (say) reagent-grade potassium hydrogen phthalate for standardizing NaOH, or pure electrolytic iron for standardizing potassium permanganate, and not think twice about it. I, however, can't get either of these things without going to a lot of trouble and probably spending a lot of money.

There are, however, substitutes, some of which are within an amateur's grasp. Sulfamic acid, for example, is available at hardware stores, and it too can be used for standardizing NaOH. Straight out of the box, though, it's probably not to be trusted, but it can be recrystallized to improve its purity.

Other compounds I have learned to make in one or two steps from commonplace chemicals. Sodium oxalate, for example, of use as a standard for KMnO₄, can be made from oxalic acid and sodium bicarbonate, both readily found. Ferrous ammonium sulfate or "Mohr's salt" can be made from ammonium sulfate and ferrous sulfate, both of which are used in gardening. I have learned to work with chemicals from grocery stores, hardware stores, photographic shops, artist's supply houses, even mineral samples. I'll never get the sort of quality that I'd get from J. T. Baker, and some of my preparations have failed or given uncertain products, but at least it's sort of fun. It's like making your own chocolate chip cookies or pretzels from scratch instead of getting them ready-made.

I'm planning to try a few volumetric experiments soon, but before I can do that, I need to prepare compounds for use in standardizing both acid-base and redox solutions. My method is to have at least two standards prepared for each reagent to be standardized and preferably three; in this way, I check the various results against each other. Some standards I already have on hand. I prepared some recrystallized potassium bitartrate for standardizing NaOH, for example, about a year ago, but the substance keeps indefinitely. Others I either don't have on hand, or have basically expired.

I intend to post notes of my work to this journal regularly and at some length and detail. It may be useful to someone else at some other time, for all I know, but mostly I'd like to have both written and photographic records of my work in one place that I can access from anywhere and which isn't going to disappear (the way that my paper notes sometimes do, although I'm striving to keep detailed and organized written notes.)

Details of particular experiments under the cut.


I've so far succeeded in preparing three standard chemicals, with more to follow.

Potassium tetroxalate dihydrate, KHC₂O₄·H₂C₂O₄·2H₂O

This is a stable double salt of potassium binoxalate and oxalic acid. It is only moderately soluble in cold water, which makes it fairly easy to crystallize. It can be useful as a standard for sodium hydroxide solutions, one mole of the tetroxalate neutralizing three moles of hydroxide. It can also be used to standardize potassium permanganate solutions. Vogel's text on quantitative analysis discourages use of potassium tetroxalate for the best work, because its water content is not entirely trustworthy, but it should be suitable for amateur work.

26 grams, about 0.206 mol, of Daly's brand oxalic acid dihydrate, was suspended in 200 mL of cold deionized water in a 500 mL beaker, and 8.3 grams of crude "pearl ash" potassium carbonate from Seattle Pottery Supply was stirred in. This is about 0.0503 mol, assuming that the "pearl ash" has a composition of K₂CO₃·³/₂H₂O. After the foam subsided, the beaker was heated to a gentle boil, driving off the excess carbon dioxide. The hot solution was filtered into a clean 500 mL beaker and cooled first to room temperature, then overnight in the refrigerator. The potassium tetroxalate began to separate upon cooling as radiating masses of small prisms (the photo is not very good.)

From Chemistry

After the overnight refrigeration the liquid was poured off the layer of crystals at the bottom of the beaker through a coarse filter. The crystalline mass was broken up with a glass rod and washed with small portions of ice-cold water. Finally the all the crystals were spread over a 12.5 cm filter paper and laid atop a pad of folded paper towels to air-dry before transfer to a clean glass bottle. The ultimate yield of potassium tetroxalate was 23.5 g or 0.0925 mol, about 92% of theoretical calculated from the amount of K₂CO₃. This value is suspiciously high.


Potassium antimonyl tartrate hemihydrate, K(SbO)C₄H₄O₆·½H₂O

This is the well-known "tartar emetic" of old-fashioned medicine, still of use in treating schistosomiasis despite its high toxicity. Some oxidizing agents used in analysis, including iodine and potassium permanganate, cleanly oxidize the Sb(III) in tartar emetic to Sb(V), and thus can be used for standardizing those oxidants, although again it is not a first choice. One reason is that tartar emetic is efflorescent, tending to lose its water of crystallization over time, and therefore its useful shelf life is limited. To make it I scaled down a recipe given in Henry Enfield Roscoe's Treatise on Chemistry (1884).

10.2 grams of potassium bitartrate, about 0.0542 mol of ordinary spice-rack cream of tartar, was placed in a 500 mL beaker along with 8.2 g or about 0.0281 mol of antimony trioxide from Seattle Pottery Supply. 100 mL of deionized water was added and the mixture slowly heated, with occasional stirring, to a gentle boil, which was kept up for about 5 minutes. By the end of this heating, there was still left a heavy white residue of undissolved antimony oxide. The very slightly yellow solution was filtered while hot into a clean Mason jar, beakers being in short supply, and cooled to room temperature. The tartar emetic began to separate in pretty, sparkling crystals of triangular and trapezoidal shape.

From Chemistry

The beaker was then refrigerated overnight, and the crystals of tartar emetic were washed with ice-cold water, collected, and air-dried in the same fashion as the potassium tetroxalate from above. The final yield was 12.2 g or about 0.0365 mol, 67% of theoretical.


Sodium tetraborate decahydrate, Na₂B₄O₇·10H₂O

This is common borax, available in any supermarket, but to be useful it must be purified. It is a fairly strong base, reacting with two moles of hydrochloric acid for every mole of tetraborate, and one of the few substances useful for direct standardization of HCl solutions. But it is efflorescent and, additionally, a bit tricky to recrystallize. From warm solutions it separates as the decahydrate, but hot solutions of borax deposit a different salt, the pentahydrate. This compound soaks up water from the air, becoming sticky, and is utterly useless as an acidimetric standard. Purification of borax requires a careful eye on temperature, which should not rise much above 50° C.

77 grams of 20-Mule-Team borax was suspended in 260 mL of deionized water in a 500 mL beaker and the probe of a programmable kitchen thermometer was placed in the mixture. This thermometer was set to sound an alarm when the temperature reached 50° C, then the beaker warmed over low heat and stirred frequently. Once the desired temperature was attained, the solution was slowly decanted through a fluted filter paper into another clean jar. Since the amount of undissolved borax left in the beaker was considerable, the solution filtered slowly, and while it drained through the funnel the beaker was kept on the burner, with care to keep the temperature between 48° C and 52° C. Once all of the solution was decanted, 25 mL more of water was added, the mixture heated to 50° C, and this final portion of liquid decanted through the filter.

From Chemistry

The filtrate was cooled at room temperature. After some time, a dusting of tiny, short prisms of borax began to separate. The jar was refrigerated overnight, as above, and the crystals washed with ice-cold water, collected, and air-dried in the same manner. The total yield of purified borax was 44.7 g, about 58% of the original amount.