Animated Solution for Chemistry - Ionic Equilibrium: Regular use of which of the following fertilizers increases the acidity of soil ?

When we think of farming and fertilizers, we often picture lush green fields and bountiful harvests. However, beneath the surface, a complex chemical balancing act is constantly taking place. The pH of the soil is a critical factor for plant growth, and the fertilizers we add can drastically alter this balance.

In this problem, we are asked to identify which common fertilizer increases the acidity of the soil over time. To solve this, we need to dive into the concept of salt hydrolysis.

Let's evaluate our options one by one. The acidity or basicity of a salt solution depends entirely on the strength of its parent acid and base.

First, we have Potassium nitrate ($KN{O}_3$). This salt is the product of a strong base, potassium hydroxide ($KOH$), and a strong acid, nitric acid ($HN{O}_3$). Because both parents are strong, neither the $K^{+}$ nor the $N{O}_3^{-}$ ions react significantly with water. The solution remains perfectly neutral.

Next is Urea ($N{H}_{2}CON{H}_2$). Urea is an organic amide. When dissolved in water, it does not immediately release $H^{+}$ or $O{H}^{-}$ ions, making it generally neutral in its direct aqueous form.

Then we have Superphosphate of lime, which contains calcium dihydrogen phosphate. While it does have some acidic protons, it is not the classic culprit for severe, long-term soil acidification compared to our final candidate.

Ammonium Sulphate

Our final option is Ammonium sulphate, $(N{H}_4{)}_{2}S{O}_4$. Let's break it down into its parent components. It is formed from the neutralization of ammonium hydroxide ($N{H}_{4}OH$), which is a weak base, and sulphuric acid ($H_{2}S{O}_4$), which is a strong acid.

When a salt of a strong acid and a weak base dissolves in water, it undergoes a process called cationic hydrolysis. Let's see how this happens in the damp soil.

First, the salt dissociates completely:

The sulphate ion ($S{O}_4^{2-}$) is the conjugate base of a strong acid, so it is incredibly weak and simply watches the reaction as a spectator. However, the ammonium ion ($N{H}_4^{+}$) is the conjugate acid of a weak base. It is highly reactive and immediately attacks the water molecules in the soil:

Notice the product of that hydrolysis reaction? It releases $H^{+}$ ions directly into the soil moisture. As the concentration of $H^{+}$ ions increases, the pH of the soil drops, making it more acidic.

Therefore, the regular application of ammonium sulphate will continuously pump $H^{+}$ ions into the ground, steadily increasing the soil's acidity. To counteract this, farmers often have to treat their fields with agricultural lime ($CaO$) to neutralize the excess acid and restore the balance.