Phosphoric acid, also known as orthophosphoric acid, is a common inorganic acid represented by the chemical formula H3PO4. It is one of the most widely produced industrial chemicals globally, often encountered as a colorless, odorless, and somewhat viscous aqueous solution. Determining the acidity of this substance requires understanding the pH scale, which measures the concentration of hydrogen ions (H+) in a solution. The pH scale ranges from 0 to 14, where values below 7 indicate acidity, above 7 indicate alkalinity, and 7 is neutral.
How Concentration Determines Measured Acidity
The pH of phosphoric acid has no single answer because its acidity is directly proportional to its concentration in water. Unlike a strong acid, which fully dissociates, phosphoric acid’s ability to release hydrogen ions depends on the amount of solvent present. Consequently, a highly concentrated solution will have a very different pH value than a highly dilute one.
Commercial-grade phosphoric acid is most frequently sold as an 85% aqueous solution, which is extremely concentrated. This viscous liquid exhibits a very low pH, typically falling around 1 or even slightly lower than 1, making it a strong irritant and corrosive substance. However, this highly acidic state changes quickly upon dilution, as the water molecules shift the equilibrium of the dissociation reactions.
A common laboratory concentration of 0.1 M phosphoric acid has a pH significantly higher, generally calculated to be in the range of 1.5 to 1.9. Further dilution to a 1% solution pushes the pH even higher, often resulting in a value between 1.5 and 2.2. This dependency demonstrates that any measurement of phosphoric acid’s acidity must be specifically tied to its molarity or percentage by weight.
The Triprotic Nature of Phosphoric Acid
Phosphoric acid’s pH is concentration-dependent due to its unique chemical structure, which classifies it as a triprotic acid. A triprotic acid possesses three acidic hydrogen atoms that it can donate sequentially in a water solution. This process occurs in three distinct, step-wise dissociation reactions rather than all at once.
Each of the three proton releases is governed by its own acid dissociation constant (pKa), with the acidity diminishing sharply after each step. The first proton is released with a pKa1 value of approximately 2.12, making it the primary contributor to the solution’s overall acidity. This relatively high pKa value, compared to values near zero for strong acids, classifies phosphoric acid as a moderately weak acid.
The second and third dissociations are substantially weaker than the first, with pKa2 around 7.21 and pKa3 near 12.32. The large gap between these pKa values shows that removing the second and third protons is chemically difficult. In typical acidic solutions, the vast majority of hydrogen ions contributing to the pH come only from the first dissociation step.
This sequential release creates a series of different phosphate species that exist in equilibrium within the solution. The chemical forms range from the fully protonated H3PO4 to the H2PO4- and HPO4(2-) ions. This complex equilibrium system enables the acid to act as an effective buffer in biological and industrial applications, as it can absorb or release hydrogen ions across a wide pH range.
Practical Uses Exploiting its Acidity
The acidic and buffering properties of phosphoric acid make it a versatile compound with numerous applications. Its capacity to lower the pH of a solution is exploited in the food and beverage industry, where it is used as an acidulant. In carbonated beverages like cola, the acid provides the characteristic tart flavor and simultaneously acts as a preservative by creating an environment too acidic for most bacteria and fungi to thrive.
In metallurgical processes, the low pH of phosphoric acid is utilized for rust removal. When applied to rusted iron or steel, the acid reacts directly with the iron oxide (rust). This reaction converts the rust into ferric phosphate, a black, inert, and more stable protective layer.
The acid also plays a role in dental procedures, primarily as an etching solution. Dentists apply a dilute form of the acid to the tooth enamel to microscopically roughen the surface before placing fillings, sealants, or orthodontic brackets. This controlled etching process creates tiny pores that allow the dental material to mechanically bond more securely to the tooth structure.