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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

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  • $\begingroup$ Iron particles can precipitated by fine filter paper and visually checked $\endgroup$ – Younus Dec 31 '18 at 9:08
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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.

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  • $\begingroup$ 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 '16 at 14:53
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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.

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