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💏intro to chemistry review

H₂PO₄⁻

Written by the Fiveable Content Team • Last updated September 2025
Written by the Fiveable Content Team • Last updated September 2025

Definition

H₂PO₄⁻ is the dihydrogen phosphate ion, a Brønsted-Lowry conjugate base that forms when a hydrogen ion (H⁺) is removed from phosphoric acid (H₃PO₄). It is an important species in acid-base equilibria and plays a key role in biological buffer systems.

5 Must Know Facts For Your Next Test

  1. H₂PO₄⁻ is the dihydrogen phosphate ion, formed when a hydrogen ion (H⁺) is removed from phosphoric acid (H₃PO₄).
  2. As a Brønsted-Lowry conjugate base, H₂PO₄⁻ can accept a hydrogen ion to form the conjugate acid, H₃PO₄.
  3. H₂PO₄⁻ is an important buffer species in biological systems, helping to maintain a stable pH in the body.
  4. The pKa of H₂PO₄⁻ is 7.2, meaning it is a moderately strong conjugate base and can effectively buffer solutions around neutral pH.
  5. H₂PO₄⁻ plays a crucial role in energy storage and transfer processes in the body, such as in the formation of ATP.

Review Questions

  • Explain the relationship between H₂PO₄⁻ and phosphoric acid (H₃PO₄) in the context of Brønsted-Lowry acid-base theory.
    • According to Brønsted-Lowry theory, H₃PO₄ is an acid that can donate a hydrogen ion (H⁺) to a base, forming the conjugate base H₂PO₄⁻. Conversely, H₂PO₄⁻ is the conjugate base of H₃PO₄, as it can accept a hydrogen ion to reform the original acid. This relationship between an acid and its conjugate base is central to understanding how H₂PO₄⁻ behaves in acid-base equilibria and buffer systems.
  • Describe the role of H₂PO₄⁻ in biological buffer systems and explain how its pKa value contributes to its buffering capacity.
    • H₂PO₄⁻ is an important buffer species in the body, helping to maintain a stable pH around neutral conditions. The pKa of H₂PO₄⁻ is 7.2, which is close to the physiological pH of 7.4. This means H₂PO₄⁻ can effectively resist changes in pH by accepting or donating hydrogen ions as needed, making it a crucial component of biological buffer systems. The ability of H₂PO₄⁻ to act as both an acid and a base allows it to neutralize both excess H⁺ and OH⁻ ions, preserving the delicate pH balance required for proper cellular function.
  • Analyze the significance of H₂PO₄⁻ in energy storage and transfer processes within the body, and explain how this relates to its role as a Brønsted-Lowry conjugate base.
    • H₂PO₄⁻ plays a crucial role in the storage and transfer of energy in the body, particularly in the form of ATP (adenosine triphosphate). As a Brønsted-Lowry conjugate base, H₂PO₄⁻ can accept a hydrogen ion to form the conjugate acid H₃PO₄, releasing energy in the process. This energy can then be used to drive various metabolic reactions, including the phosphorylation of ADP to form ATP. The reversible nature of this reaction, with H₂PO₄⁻ acting as both an acid and a base, allows for the efficient storage and release of energy as needed by the body's cells. This highlights the important functional role of H₂PO₄⁻ in the context of Brønsted-Lowry acid-base theory and its broader significance in cellular bioenergetics.