I have a rusty metal rainwater tank, and I'd like to find the amount of iron in the water. (The water I've collected is perfectly clear.)

How would chemists in the 1800's have gone about this? What chemicals/apparatus would I need to purchase?

So far, I've soaked a tissue in the water and am letting it air dry to see if any rust discoloration becomes apparent.

The following info from the WHO might be helpful:

Iron (as Fe2+) concentrations of 40 µg/litre can be detected by taste in distilled water. In a mineralized spring water with a total dissolved solids content of 500 mg/litre, the taste threshold value was 0.12 mg/litre. In well-water, iron concentrations below 0.3 mg/litre were characterized as unnoticeable, whereas levels of 0.3–3 mg/litre were found acceptable (E. Dahi, personal communication, 1991).

In drinking-water supplies, iron(II) salts are unstable and are precipitated as insoluble iron(III) hydroxide, which settles out as a rust-coloured silt. Anaerobic groundwaters may contain iron(II) at concentrations of up to several milligrams per litre without discoloration or turbidity in the water when directly pumped from a well, although turbidity and colour may develop in piped systems at iron levels above 0.05–0.1 mg/litre. Staining of laundry and plumbing may occur at concentrations above 0.3 mg/litre (4).

Iron also promotes undesirable bacterial growth ("iron bacteria") within a waterworks and distribution system, resulting in the deposition of a slimy coating on the piping (4). http://www.who.int/water_sanitation_health/dwq/chemicals/iron.pdf

  • $\begingroup$ Iron particles can precipitated by fine filter paper and visually checked $\endgroup$
    – Younus
    Dec 31, 2018 at 9:08

3 Answers 3


In a rainwater tank in contact with air, dissolved iron is probably oxidized to iron(III) ($\ce{Fe^3+}$).

A simple and sensitive test for $\ce{Fe^3+}$ in water uses thiocyanate ions ($\ce{SCN-}$, also known as rhodanide), which form the blood-red coloured complexes $\ce{[Fe(SCN)(H2O)5]^2+}$, $\ce{[Fe(SCN)2(H2O)4]+}$, and $\ce{[Fe(SCN)3(H2O)3]}$.

$$\ce{[Fe(H2O)6]^3+ + SCN- <=> [Fe(SCN)(H2O)5]^2+ + H2O}$$

Take a precise volume of the water (e.g. $10\ \mathrm{ml}$), acidify with dilute sulfuric acid (e.g. $2\ \mathrm{ml}$), add a thiocyanate solution (e.g. $5\ \mathrm{ml}$ $\ce{KSCN}$ or $\ce{NH4SCN}$, $c =0.5{–}2\ \mathrm{mol/l}$), and fill up with distilled water to a precise volume (e.g. $25\ \mathrm{ml}$).

An orange to red colour should appear.

The colour of the solution can be compared to that of standard solutions that contain known concentrations of $\ce{Fe^3+}$, e.g. prepared from iron(III) nitrate ($\ce{Fe(NO3)3.9H2O}$).

In a laboratory, the intensity of the red colour could be measured using a photometer.

Similar simple test kits are commercially available, e.g. for aquariums, garden ponds, or other water or soil samples.

  • $\begingroup$ I agree with colorimetry being the method of choice. It cannot be very precise by definition, but you just cannot make anything wrong. It is very robust and does not need precise scale. The standards can be prepared by further dilution in known ratios. You could also use your camera as a very imprecise photometer. $\endgroup$
    – ssavec
    Feb 24, 2016 at 14:53
  • $\begingroup$ My stock method is to use 2,2'-bipy and iron to make the red complex for measurement, my method would require you to add a reducing agent such as hydroxylamine to get all the dissolved iron into the form of iron(II). But the thiocyanate method would work, $\endgroup$ Oct 16, 2022 at 7:36
  • $\begingroup$ I would disagree with Ssavec, when used with care colorimetry can be used to get very precise results $\endgroup$ Oct 16, 2022 at 9:21

If you want to test for the presence of iron(II) in solution, you could add $\ce{CO3^{2-}}$ ions to form a green precipitate. This could be extracted and measured, although it could be a very small amount unless you concentrated the sample water by evaporating it down. The main tool would be an accurate balance, but also some liquid measuring equipment.

An alternative method that would be more analytically accurate would be a redox titration using something like potassium permanganate or dichromate. This would require an accurately known concentration of the reagent, and some more specific equipment like a burette and a stand, and a volumetric flask for precise liquid measurements.

  • $\begingroup$ Very small amounts of $\ce{Fe^{2+}}$ ions are quite improbable in water, because they are quickly oxidized by atmospheric oxygen, and transformed into $\ce{Fe^{3+}}$ ions. And this $\ce{Fe^{3+}}$ ion does not make a green precipitate with $\ce{CO3^{2-}}$ ions. It makes a brown precipitate. The same critics could be made for the second method, using potassium permanganate, which only reacts with $\ce{Fe^{2+}}$ ion and not with $\ce{Fe^{3+}}$ ions $\endgroup$
    – Maurice
    Oct 16, 2022 at 9:49

There is a problem which I think that the other people might not be aware of. It is a problem which is well known in nuclear chemistry. If you take a sample of lake water from some random lake such as a Swedish lake then if you measure the plutonium content of the water by digesting the sample and then filtering then you get a higher value for fallout plutonium than if you filter it first and then digest.

While iron is very different to plutonium, I think that the same problem could exist. The amount of iron which is present as a fine (colidal solid) might be important. I would suggest the following.

Take a large sample of water from the tank, now filter part of this through the finest filter you can find. The other part you should combine with a known volume of hydrochloric acid and boil it up.

After cooling the boiled acidic water should be filtered. I would take both filtrates and then measure the iron by means of colourmetric tests.

My normal test for iron uses 2,2'-bipy to form a complex with the iron(II), this test can be used to determine what fraction of the iron is present as iron(II).

I would make up four volumetric flasks with the filtered iron solutions, I would put in each flask ammonium aceate (to buffer the mixture), 2,2'-bipy solution and the iron sample. I have found that this method works down to ppm levels of iron. To two of the flasks I would add hydroxylamine solution to reduce all the iron to iron(II). I would also make up a control (reagent blank) flask and a set of standards in the range 0 to 10 ppm of iron. I would then use a UV / vis machine to measure the red complex [Fe(bipy)3]2+

I think that a late victorian chemist would be able to make 2,2'-bipy from pyridine, but if you want a method which an early victorian could have used then use potassium thiocyanate to form the red complex with the iron(III).

For that method you would need to put in the flasks potassium thiocyanate, you might want to read about the Fricke dosimeter which is used in nuclear chemistry to measure very high radiation doses. It is based on the conversion of iron(II) in acidic media which is oxygen saturated. The conversion to form iron(III) is under some conditions proportional to the radiation dose. Some people use thiocyanate in the iron(III) determination to make it more senstive.

I suggest you read

The International Journal of Applied Radiation and Isotopes Volume 29, Issue 3, March 1978, Pages 151-157 https://doi.org/10.1016/0020-708X(78)90137-0

If you want to dig deeply into the problem of how to measure iron(III) with thiocyanate


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