# Does adding slices of lemons to water reduce its chlorine content?

Often restaurants add slices of lemon to water. They seem to make the water taste better and perhaps reduce or at least cover up the chlorine taste.

I've heard that Vitamin C tablets can be used to remove chlorine from tap water.

Q: Does anyone know if lemons, being rich in vit C, will encourage the removal of chlorine from tap water?

• May 12 '17 at 16:32
• My only guess would be that adding an organic compound will allow the residual Cl to react with and oxidize the compound, thus reducing the Cl content. May 15 '17 at 18:00

Yes , a quick and cheap way to reduce chlorine residual is to add ascorbic acid (VITAMIN C), which if added in an equal molar amount to free chlorine would completely eliminate chlorine residual within 1 min.

I am quoting from ABSTRACT from a paper from, Int. Journal of Engineering Research and Applications , ISSN : 2248-9622, Vol. 3, Issue 5, Sep-Oct 2013, pp.1647-1651.

"Society’s preference for bottled water mainly comes from a dislike of the mildly pungent taste and odor of tap water. Such odor and taste may be caused by chlorine residual (in the form of free chlorine or chloramines), which is maintained in a water distribution system to prevent regrowth of microorganisms and is thus present in tap water.

"Two methods of reducing chlorine residual, boiling tap water and adding ascorbic acid to tap water, were tested. The results showed that, for free chlorine, a quick and cheap way to reduce chlorine residual is to add ascorbic acid, which if added in an equal molar amount to free chlorine would complete eliminate chlorine residual within 1 min "

• Reduction of Chlorine Residual in Tap Water with Temperature

As shown aboce, the chlorine residual decreased faster at higher temperatures.

• Reduction of Free Chlorine and Monochloramine in Tap Water through Boiling

As shown in Figure 2, both residuals of free chlorine and monochloramine decreased with boiling time, and the residual of free chlorine decreased slightly faster than that of monochloramine.

The boiling time required to completely remove free chlorine was about 33 min.$$\ce{^1}$$

• Effect of Ascorbic Acid on Free Chlorine and Monochloramine

Ascorbic acid added in a molar ratio of 1:1 (free chlorine to ascorbic acid) completely eliminated free chlorine. The DPD indicator(A DPD solution can also be used for the measurement of chlorine residual) did not produce a red color when added within 1 min,indicating complete elemination chlorine residual within 1 min.$$\ce{^2}$$

The reduction of free chlorine by ascorbic acid can be expressed as follows:

$$\ce{C6H8O6}$$ + $$\ce{HOCl}$$$$\ce{C6H6O6}$$ +$$\ce{HCl}$$ + $$\ce{H2O}$$

Ascorbic acid $$\ce{(C6H8O6)}$$ is oxidized to dehydroascorbic acid $$\ce{(C6H6O6)}$$ . Simultaneously, free chlorine $$\ce{(HOCl)}$$ is reduced to chloride ions (HCl).

In the experiment conducted it was determined that 10 mg of ascorbic acid is required to treat 1 L of tap water. and the quickest way to reduce the concentration of free chlorine in tap water is to add powdered vitamin C (ascorbic acid) tablets in less than 1 min.$$\ce{^3}$$

Thanks to N4v i will be using this cost effective method to remove chlorine from water.(wow).

Reference

Yes, vitamin C can be used to remove dissolved $$\ce{Cl_2}$$ from water. The same applies for dissolved $$\ce{Br_2}$$ and $$\ce{I_2}$$. Whether a lemon slice (or several) has enough vitamin C to quantitatively react with all the dissolved $$\ce{Cl_2}$$ in a glass of water is another matter and also a simple hypothesis that can be tested easily in an undergraduate quantitative analysis or intrumental analysis lab. The next two paragraphs are from my old pre-lab lab write-up where students used in situ generated iodine for the quantitative determination of vitamin C in commercial 1 g vitamin C tablets. Bromine, generated in situ by oxidizing bromide ions, also works well. We never tried generating $$\ce{Cl_2}$$ by in situ oxidation of chloride ions because of the interfering side oxidation reaction.

Commercial vitamin C tablets are not pure vitamin C: it is not suitable for formulation in pure form. Rather, commercial vitamin C tablets are comprised of pure vitamin C (ascorbic acid, denoted as HAsc) plus various binders, processing additives and excipients. To determine how much HAsc is in a tablet, we determine the number of moles of HAsc in the tablet and use the molar mass of HAsc. From the measured mass of the tablet, the % purity is then computed. Alternatively, we may determine the % purity of the tablet and then use its measured mass to compute the mass of pure HAsc the tablet contained. Either way, it is necessary to determine the number of moles of HAsc in the tablet or in a measured fraction of the tablet. In practice, the latter is done since commercial 1 g vitamin C tablets contain inconveniently large amounts of HAsc.

The technique of coulometric titration is used to determine the number of moles of HAsc in the titrated aliquoted of solution prepared from the tablet. Iodine is capable of rapidly and irreversibly oxidizing HAsc to dehydroacsorbic acid, with iodide ion being another product. Molecular iodine is poorly soluble in aqueous solution, unless iodide ion is present, in which case the solubility is much higher due to formation of the tri-iodide ion. The tri-iodide ion rapidly and readily releases iodine molecules for reaction with HAsc in aqueous solution. Thus, since one HAsc molecule reacts with one diatomic iodine molecule, the iodine molecule can be generated electrolytically in solution by oxidizing a pair of iodide ions, which are deliberately present in excess in the titration vessel. Therefore, removing 2 electrons fron 2 iodide ions yields 1 iodine diatomic molecule which then reacts quantitatively with one HAsc molecule. The 2 iodide ions are thus regenerated and keep cycling around. The stoichiometry is 2 equivalents per mole of HAsc. (One equivalent is the number of moles of electrons required to reduce one mole of hydrogen ions to hydrogen atoms.)