Chemical identity and nomenclature
Monopotassium phosphate is an ionic compound commonly written as KH2PO4. The same substance is widely called potassium dihydrogen phosphate. In water it supplies potassium ions, K+, and the dihydrogen phosphate ion, H2PO4−, which belongs to the triprotic phosphoric acid system.
The formula KH2PO4 corresponds to one potassium per phosphate unit with two acidic hydrogens retained. Related salts include K2HPO4 (dipotassium phosphate) and K3PO4 (tripotassium phosphate), which differ in charge state and acid–base behavior.
Composition and molar mass
The molar mass follows directly from atomic composition. Using standard atomic masses (to typical classroom precision):
| Element | Count | Contribution (g/mol) |
|---|---|---|
| K | 1 | \(\approx 39.10\) |
| H | 2 | \(\approx 2 \times 1.008 = 2.016\) |
| P | 1 | \(\approx 30.97\) |
| O | 4 | \(\approx 4 \times 16.00 = 64.00\) |
Summation:
\[ M(\mathrm{KH_2PO_4}) \approx 39.10 + 2.016 + 30.97 + 64.00 = 136.086\ \mathrm{g/mol} \approx 136.09\ \mathrm{g/mol} \]
Dissociation in water and electrolyte character
As an ionic salt, monopotassium phosphate dissociates extensively in water:
\[ \mathrm{KH_2PO_4(s) \rightarrow K^+(aq) + H_2PO_4^-(aq)} \]
Potassium, K+, is the spectator cation of a strong base (KOH) and has negligible acid–base influence. The pH behavior is dominated by H2PO4−.
Acid–base behavior of the dihydrogen phosphate ion
The dihydrogen phosphate ion is amphiprotic: it can donate a proton to form HPO42− and can accept a proton to form H3PO4. These equilibria connect to the successive dissociations of phosphoric acid:
\[ \mathrm{H_3PO_4 \rightleftharpoons H^+ + H_2PO_4^-} \quad (K_{a1},\ pK_{a1}\approx 2.15) \] \[ \mathrm{H_2PO_4^- \rightleftharpoons H^+ + HPO_4^{2-}} \quad (K_{a2},\ pK_{a2}\approx 7.20) \] \[ \mathrm{HPO_4^{2-} \rightleftharpoons H^+ + PO_4^{3-}} \quad (K_{a3},\ pK_{a3}\approx 12.38) \]
A useful approximation for an aqueous solution containing primarily an amphiprotic species such as H2PO4− is:
\[ \mathrm{pH \approx \tfrac{1}{2}\left(pK_{a1}+pK_{a2}\right)} \approx \tfrac{1}{2}(2.15+7.20)=4.675 \]
The mild acidity reflects the balance between the two tendencies of H2PO4−. Proton donation governed by \(K_{a2}\) is typically stronger than proton acceptance governed by \(K_{b} = K_w/K_{a1}\), shifting equilibrium toward pH below 7 for typical concentrations.
Buffer relevance in the phosphate system
Monopotassium phosphate is frequently paired with dipotassium phosphate, K2HPO4, to create a phosphate buffer near neutral pH. The conjugate pair is H2PO4−/HPO42−, and the Henderson–Hasselbalch relation applies:
\[ \mathrm{pH = pK_{a2} + \log\!\left(\frac{[HPO_4^{2-}]}{[H_2PO_4^-]}\right)} \]
When \([HPO_4^{2-}] = [H_2PO_4^-]\), the logarithmic term is zero and \(\mathrm{pH \approx p}K_{a2}\), placing the buffer region near \(\mathrm{pH \approx 7.2}\).
Common confusions and practical notes
- Monopotassium phosphate (KH2PO4) versus dipotassium phosphate (K2HPO4): different conjugate-base forms and different typical pH ranges in water.
- “Mono” referring to potassium count, not to the number of acidic hydrogens in the phosphate framework.
- Phosphate buffering centered near \(pK_{a2}\) rather than near \(pK_{a1}\), reflecting the H2PO4−/HPO42− pair.