Dipotassium phosphate (\(\text{K}_2\text{HPO}_4\)) is a common substance used in food, fertilizers, and laboratories. When dissolved in water, this compound forms a solution that is slightly basic (pH above 7). This basicity is determined by the chemical behavior of the hydrogen phosphate ion (\(\text{HPO}_4^{2-}\)), which is released when the salt dissociates. The overall acidic or basic character results from a competition of chemical reactions the ion undergoes with water.
Understanding Acids, Bases, and Salts
To understand \(\text{K}_2\text{HPO}_4\), it is necessary to define acids, bases, and salts. Under the Brønsted-Lowry definition, an acid donates a proton (\(\text{H}^+\)), while a base accepts one.
The pH scale measures acidity, ranging from 0 to 14. A pH below 7 is acidic, above 7 is basic, and 7 is neutral. Salts are ionic compounds formed from the reaction of an acid and a base, which dissolve in water to produce ions.
When a salt dissolves, its component ions may react with water (hydrolysis), changing the solution’s pH. Ions derived from weak acids or weak bases can accept or donate protons to water, influencing the final pH.
The Chemistry of the Hydrogen Phosphate Ion
The \(\text{HPO}_4^{2-}\) ion originates from phosphoric acid (\(\text{H}_3\text{PO}_4\)), which is a triprotic acid capable of donating three protons in a stepwise manner.
The first deprotonation step forms the dihydrogen phosphate ion (\(\text{H}_2\text{PO}_4^-\)). The second step involves \(\text{H}_2\text{PO}_4^-\) losing a proton to form the hydrogen phosphate ion (\(\text{HPO}_4^{2-}\)), the ion of interest in \(\text{K}_2\text{HPO}_4\). The third proton can be removed from \(\text{HPO}_4^{2-}\) to form the fully deprotonated phosphate ion (\(\text{PO}_4^{3-}\)). The hydrogen phosphate ion is thus an intermediate species within the phosphoric acid system.
Amphiprotic Behavior
The hydrogen phosphate ion (\(\text{HPO}_4^{2-}\)), which dissociates from \(\text{K}_2\text{HPO}_4\), is classified as an amphiprotic species, meaning it can act as both an acid and a base. This dual nature exists because it possesses a proton it can donate and also carries a negative charge allowing it to accept a proton.
When \(\text{HPO}_4^{2-}\) is in water, two competing reactions occur simultaneously. Acting as a weak acid, it donates its proton to water, producing the phosphate ion (\(\text{PO}_4^{3-}\)) and hydronium ion (\(\text{H}_3\text{O}^+\)), which contributes to acidity.
Acting as a weak base, it accepts a proton from water, forming the dihydrogen phosphate ion (\(\text{H}_2\text{PO}_4^-\)) and a hydroxide ion (\(\text{OH}^-\)), which contributes to basicity. The final pH of the solution depends on which of these opposing reactions predominates.
Calculating the Predominant Character
To determine if the solution is acidic or basic, the relative strengths of the ion’s acidic and basic tendencies are compared using their equilibrium constants: the acid dissociation constant (\(\text{K}_a\)) and the base dissociation constant (\(\text{K}_b\)).
The \(\text{K}_a\) value for \(\text{HPO}_4^{2-}\) acting as an acid is the third dissociation constant of phosphoric acid (\(\text{K}_{a3}\)), approximately \(4.2 \times 10^{-13}\).
The \(\text{K}_b\) value for \(\text{HPO}_4^{2-}\) acting as a base is calculated using the relationship \(\text{K}_w = \text{K}_a \times \text{K}_b\), where \(\text{K}_w\) is \(1.0 \times 10^{-14}\). Since \(\text{HPO}_4^{2-}\) is the conjugate base of \(\text{H}_2\text{PO}_4^-\), its \(\text{K}_b\) is derived from the second dissociation constant (\(\text{K}_{a2} \approx 6.2 \times 10^{-8}\)).
The calculated \(\text{K}_b\) is approximately \(1.6 \times 10^{-7}\). Comparing the constants, \(\text{K}_b\) (\(1.6 \times 10^{-7}\)) is significantly larger than \(\text{K}_a\) (\(4.2 \times 10^{-13}\)). Therefore, the basic reaction, where \(\text{HPO}_4^{2-}\) accepts a proton to form \(\text{OH}^-\), predominates, resulting in a slightly basic solution with a pH typically around 8.
Real-World Uses of Dipotassium Phosphate
Dipotassium phosphate is a versatile compound valued for its solubility and mild basicity. It has several key applications:
- Food additive: It functions as a stabilizer and emulsifier, particularly in non-dairy creamers to prevent coagulation. Its basic nature also helps regulate food acidity, and it serves as a source of potassium.
- Agriculture: It is used as a fertilizer to supply essential potassium and phosphorus nutrients for plant growth.
- Buffer solutions: In research, it is frequently used alongside monopotassium phosphate (\(\text{KH}_2\text{PO}_4\)) to maintain a stable, near-neutral pH crucial for biological processes.
- Medical supplement: It is used to replenish the body’s phosphate levels and help maintain acid-base balance.