Sodium phosphate, chemically known as \(\text{Na}_3\text{PO}_4\) or Trisodium Phosphate (TSP), is a white, crystalline substance that is highly soluble in water. This compound readily dissolves, meaning a significant amount of the solid can be incorporated into an aqueous solution. At \(25^\circ\text{C}\), the solubility of \(\text{Na}_3\text{PO}_4\) is approximately \(14.5\) grams per \(100\) milliliters of water. This high solubility makes it useful in various applications, including as a cleaning agent and a food additive.
The Fundamental Rule of Solubility
The high solubility of sodium phosphate is determined by fundamental rules of ionic chemistry, specifically concerning the cation it contains. Sodium phosphate is an ionic compound composed of the positively charged sodium cation (\(\text{Na}^+\)) and the negatively charged phosphate polyatomic anion (\(\text{PO}_4^{3-}\)).
The most dominant principle governing this compound’s solubility is the universal rule for alkali metals. This rule states that all compounds containing alkali metal ions, such as lithium, potassium, and sodium, are soluble in water. The presence of three sodium cations in the chemical formula overrides other factors that might typically make the phosphate portion of the molecule insoluble.
Most phosphate-containing compounds, like calcium phosphate, are generally insoluble in water. However, the strong tendency of the sodium ion to dissolve ensures that sodium phosphate is completely soluble. If a compound contains an alkali metal ion, its solubility is practically guaranteed, regardless of the accompanying anion.
The Dissolution Process: Complete Dissociation
When sodium phosphate is added to water, it undergoes a process called dissociation, separating completely into its constituent ions. This breaking apart of the ionic lattice is a defining characteristic of highly soluble salts.
The process is represented by the chemical equation: \(\text{Na}_3\text{PO}_4(s) \rightarrow 3\text{Na}^+(aq) + \text{PO}_4^{3-}(aq)\). Polar water molecules surround and stabilize the separated ions. This stabilization, known as solvation, prevents the positive and negative ions from reattracting and reforming the solid.
The highly polar nature of water effectively pulls the individual sodium cations and phosphate anions away from the crystal structure. Because the compound dissociates completely into charged ions, the resulting solution is an excellent conductor of electricity, classifying sodium phosphate as a strong electrolyte.
Resulting Solution Chemistry: Why Sodium Phosphate is Basic
The complete dissociation of sodium phosphate leads to a solution that is distinctly alkaline, meaning it has a high pH. This basicity is not caused by the sodium ion, which is considered a spectator ion because it does not react further with the water. The chemical property of the solution is instead driven by the phosphate anion (\(\text{PO}_4^{3-}\)).
The phosphate ion is the conjugate base of a weak acid, phosphoric acid, meaning the phosphate ion itself is a reasonably strong base. When the phosphate ion is released into the water, it reacts with water molecules in a process called hydrolysis. During hydrolysis, the phosphate ion pulls a hydrogen ion (\(\text{H}^+\)) from a water molecule.
This reaction generates hydroxide ions (\(\text{OH}^-\)) in the solution. The hydrolysis reaction is represented as \(\text{PO}_4^{3-}(aq) + \text{H}_2\text{O}(l) \rightleftharpoons \text{HPO}_4^{2-}(aq) + \text{OH}^-(aq)\). The accumulation of excess hydroxide ions causes the pH of a sodium phosphate solution to rise significantly; a \(1\%\) solution can reach a pH of \(12\). This strong alkalinity is why Trisodium Phosphate is highly effective as a powerful cleaning and degreasing agent.