Animated Solution for Chemistry - Environmental Chemistry: A water sample has ppm level concentration of following anions : $F^{-}=10$ $S{O}_4^{2-}=100$ $N{O}_3^{-}=50$ The anion/anions that make/makes the water sample unsuitable for drinking is/are

When you pour a glass of water, its crystal-clear appearance might give you a false sense of security. But true water quality is not determined by what you can see; it is dictated by the invisible dissolved ions swimming within it.

Environmental chemistry plays a crucial role in establishing the safe limits for these ions. If the concentration of certain anions exceeds the maximum prescribed limits, the water transforms from a life-sustaining liquid into a slow poison.

In this problem, we are tasked with evaluating a water sample containing three specific anions: Nitrate ($N{O}_3^{-}$), Sulphate ($S{O}_4^{2-}$), and Fluoride ($F^{-}$). Let's put each of these suspects under the microscope.

The Fertilizer Runoff

Our first suspect is the Nitrate ion ($N{O}_3^{-}$). The given water sample contains $50\text{ ppm}$ of nitrate.

Where does this nitrate come from? The primary culprit is agricultural runoff. Farmers use nitrogen-rich fertilizers to boost crop yields. When it rains, these highly soluble nitrates wash into the soil and eventually infiltrate our groundwater reservoirs.

According to international drinking water standards, the maximum permissible limit for nitrate is exactly $50\text{ ppm}$. Since our sample is right at the threshold, it is technically still considered suitable for drinking.

But what happens if this limit is breached? Inside the human body, particularly in infants, nitrate is reduced to nitrite ($N{O}_2^{-}$) by bacteria in the stomach. This nitrite enters the bloodstream and oxidizes the iron in hemoglobin from its ferrous state ($F{e}^{2+}$) to its ferric state ($F{e}^{3+}$).

This altered molecule is called methemoglobin, and it is completely incapable of binding and transporting oxygen. The resulting oxygen starvation gives the infant's skin a bluish tint, a deadly condition known as Methemoglobinemia or "Blue Baby Syndrome".

The Industrial Footprint

Next, we examine the Sulphate ion ($S{O}_4^{2-}$). Our water sample contains $100\text{ ppm}$ of sulphate.

Sulphates find their way into water bodies through a variety of channels. Acid rain, which contains dilute sulfuric acid, is a major natural source. On the anthropogenic side, industrial effluents from mining and chemical manufacturing heavily contribute to sulphate pollution.

Fortunately, the human body is quite tolerant to sulphates. The maximum prescribed limit for sulphate in drinking water is a generous $500\text{ ppm}$.

Since our sample only contains $100\text{ ppm}$, it is well within the safe zone. However, if you were to drink water exceeding the $500\text{ ppm}$ limit, the excess sulphate would draw water into your intestines through osmosis, leading to a severe laxative effect and dehydration.

The Double-Edged Sword

Finally, we arrive at the Fluoride ion ($F^{-}$). The sample contains $10\text{ ppm}$ of fluoride.

Fluoride is a fascinating ion because its health effects are highly concentration-dependent. It is a classic double-edged sword. At low concentrations (around $1\text{ ppm}$), fluoride is actually beneficial. It reacts with the hydroxyapatite in our tooth enamel to form fluorapatite:

Fluorapatite is significantly harder and more resistant to the acids produced by bacteria in our mouths, thereby preventing tooth decay.

However, the maximum permissible limit for fluoride is strictly capped at $1.5\text{ ppm}$.

Our sample contains a staggering $10\text{ ppm}$! When fluoride concentrations exceed $2\text{ ppm}$, it causes brown mottling of the teeth. But at levels above $10\text{ ppm}$, the consequences become catastrophic. The excess fluoride begins to interfere with calcium metabolism, leading to skeletal fluorosis—a painful condition characterized by severe bone deformities and joint damage.

We have interrogated all three suspects. The nitrate is at the borderline but safe. The sulphate is well within the acceptable range.

But the fluoride concentration is dangerously high, far exceeding the $1.5\text{ ppm}$ limit. Therefore, it is the Fluoride ($F^{-}$) alone that renders this water sample completely unsuitable for drinking.

This problem is a powerful reminder that environmental chemistry is not just about memorizing numbers; it is about understanding the delicate chemical balance that sustains human health.