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We know that iron(II) ions in water have a green colour. Why we cannot draw the spectrum of iron(II) directly in distilled water to measure the absorbance then determine the concentration of iron(II)? We generally use ortho-phenanthroline as a complex agent to form a red-orange coloured complex with iron(II), then measuring the absorbance and the concentration. Why we Cannot use just distilled water as complex agent knowing that it gives green colour with iron(II)?

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    $\begingroup$ The complex is much more strongly light absorbing than the iron ions themselves. So you are putting a bell on the cat: much easier to hear when you startle the cat with light! $\endgroup$
    – Ed V
    Mar 20, 2021 at 17:58
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    $\begingroup$ Iron (II) has no color to speak of. You can't see it until you have a lot of it. $\endgroup$
    – Ivan Neretin
    Mar 20, 2021 at 18:03
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    $\begingroup$ The last photo here show how faint the green color was. So the idea is to have a strongly colored complex and have the target species and complexing agent be colorless or only lightly colored. $\endgroup$
    – Ed V
    Mar 20, 2021 at 18:45

2 Answers 2

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I'm sure OP get the answer by reading all comments and the answer elsewhere. However, I'd like to point out few things on this spectrophotometric determination of $\ce{Fe^2+}$ in drinking water (it is not distilled water):

  1. Iron can exist in natural waters where it originates from the dissolution of minerals. Even though most iron containing minerals are not very soluble in water as evident by the solubility products of iron-containing salts. However, a trace amount of iron can still dissolve in water. The dissolve iron is predominantly found in the form of $\ce{Fe^3+}$, because dissolved oxygen in water oxidizes $\ce{Fe^2+ -> Fe^3+}$.
  2. Based on above fact, water cant be green, but probably yellow. But, it can be still determined by UV-vis spectroscopy, only if you have a lot of $\ce{Fe^3+}$ ions in your unknown drinking water sample to do so (see Ivan Neretin's comment and Maurice's answer), which is not possible.
  3. On the other hand, 3 molecules of 1,10-phenanthroline (ortho-phenanthroline, a bidentate) chelate with one ion of $\ce{Fe^2+}$ to give highly intense orange-red $\ce{[Fe(o-phen)3]^2+}$ complex ion in water (Ref.1). The equilibrium constant for this reaction $\left(\ce{Fe^2+ + 3 o-phen -> [Fe(o-phen)3]^2+}\right)$ is $2.5 \times 10^6$ at $\pu{25 ^\circ C}$ in the $\mathrm{pH}$ range between $\mathrm{pH} = 3$ and $\mathrm{pH} = 9$. Thus, a $\mathrm{pH}$ of about $3.5$ (using acetate buffer) is recommended to prevent precipitation of iron salts.
  4. The molar absorption coefficient (ε) of this ferrous complex, $\ce{[Fe(o-phen)3]^2+}$ is $\pu{11,100 dm3 mol-1 cm-1}$ at the wavelength of maximum absorbance, $\lambda_{max} = \pu{508 nm}$. Thus, spectrophotometric absorbance of this complex can be easily measured in $\pu{10^{-4} M}$ to $\pu{10^{-5} M}$ or even lower levels at $\pu{508 nm}$, which is compatible with $[\ce{Fe^3+}]$ levels in drinking water (usually well water).
  5. However, there is a problem: 1,10-Phenanthroline complexes quantitatively only with $\ce{Fe^2+}$ ions. But that can be solve using a reducing agent such as hydroxyl amine ($\ce{HCl}$ salt, which is soluble in water): $$\ce{4Fe^3+ + 2HO-NH2.HCl -> 4Fe^2+ + N2O + 6H+ + 2Cl- + H2O}.$$ The reducing agent should be used in excess to prevent $\ce{Fe^2+}$ oxidizes back to $\ce{Fe^3+}$ by dissolve oxygen.

References:

  1. R. Belcher, "The application of chelate compounds in analytical chemistry," Pure Appl. Chem. 1973, 34(1), 13-28 (DOI: http://dx.doi.org/10.1351/pac197334010013).
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    $\begingroup$ You forgot HCl in the right hand side of the last equation. $\endgroup$
    – Maurice
    Mar 21, 2021 at 9:45
  • $\begingroup$ @: Thanks for the notation. I corrected it. $\endgroup$
    – Mathew Mahindaratne
    Mar 21, 2021 at 15:14
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The o-phenanthrolin is extraordinarily sensitive to the smallest amount of $\ce{Fe^{2+}}$. A concentration as low as of $2$ microgramms $\ce{Fe^{2+}}$ per liter is still colored and its concentration can be measurable by optical absorption. As a comparison, a concentration of $\pu{30g/L}$ $\ce{Fe^{2+}}$ is the minimum measurable by Beer's law in water.

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    $\begingroup$ The answer would be greatly improved if you gave the detection limits for just an aqueous solution of $\ce{Fe^{2+}}$. Or perhaps you could compare the molar absorptivities. $\endgroup$
    – MaxW
    Mar 20, 2021 at 19:36

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