Polyprotic acids are acids that can donate more than one proton, H+, per molecule. Examples include carbonic acid, H2CO3, sulfuric acid, H2SO4, and phosphoric acid, H3PO4. They matter because each proton is released in a separate step, so their solutions can contain several related chemical species at once.
This stepwise behavior controls pH, buffering, titration curves, and many biological and environmental systems.
Each ionization step has its own acid dissociation constant, such as Ka1, Ka2, and Ka3. Usually Ka1 is larger than Ka2, and Ka2 is larger than Ka3, because it becomes harder to remove a positive proton from an increasingly negative ion. At different pH values, different forms of the acid dominate, such as H3A, H2A-, HA2-, or A3-.
During titration with a strong base, a polyprotic acid can produce multiple buffer regions and multiple equivalence points if the Ka values are far enough apart.
Understanding Chemistry: Polyprotic Acids
When the first proton leaves an acid molecule, water usually accepts it and becomes hydronium. What remains is not simply an inactive leftover. It is a new particle with its own acid base behavior.
It can accept a proton back, or it can lose another proton. The charge on this particle strongly affects the next step. Removing a proton from a neutral molecule is often much easier than removing one from a negatively charged ion.
The negative ion attracts the positive proton and resists becoming even more negative. This electrostatic effect explains the usual pattern of weakening acid strength through successive ionizations. There are exceptions in detail, especially when molecular structure stabilizes a charge in an unusual way.
A solution of a polyprotic acid is best viewed as a mixture whose composition changes smoothly with pH. Near a particular acid dissociation value, two neighboring forms are both important. At the pH equal to that value, those two forms have equal amounts.
A pH one unit below that value gives roughly ten times more of the protonated form. A pH one unit above it gives roughly ten times more of the deprotonated form. This pattern helps students predict which form dominates without solving every equilibrium calculation.
The middle forms are especially interesting because they can donate a proton or accept one. Such particles are called amphiprotic. Their behavior depends on the surrounding solution.
Titration makes the separate proton removals visible through stoichiometry. Each mole of a strong base added can remove one mole of acidic protons. For a substance with three removable protons, complete neutralization needs three moles of hydroxide for every mole of acid.
The first equivalence point occurs after the first proton has been neutralized. At that point, the solution mainly contains the form that has lost one proton. Later equivalence points follow as more base is added.
In a real graph, separate jumps only appear clearly when the acid dissociation values are far apart. If they are close together, the changes overlap and the graph may look like it has fewer distinct stages. Indicator choice must match the pH range near the endpoint.
Carbonic acid shows why these ideas matter outside a laboratory. It is part of the carbon dioxide system in blood, lakes, oceans, and fizzy drinks. Its related ions help resist sudden pH changes.
Phosphate systems perform a similar buffering role inside cells and in soil. In water treatment, carbonate chemistry affects hardness and scale formation in pipes. When solving problems, first identify the relevant proton loss step.
Do not treat every proton as equally easy to remove. Check whether a stated concentration is large enough that later steps contribute meaningfully to the hydrogen ion concentration. In titration questions, keep equilibrium reasoning separate from mole counting until the stoichiometric reaction is complete.
Key Facts
- A polyprotic acid can donate two or more protons in separate ionization steps.
- For a triprotic acid: H3A ⇌ H+ + H2A-, Ka1 = [H+][H2A-]/[H3A].
- Second ionization: H2A- ⇌ H+ + HA2-, Ka2 = [H+][HA2-]/[H2A-].
- Third ionization: HA2- ⇌ H+ + A3-, Ka3 = [H+][A3-]/[HA2-].
- For most polyprotic acids, Ka1 > Ka2 > Ka3, so pKa1 < pKa2 < pKa3.
- At pH = pKa for a step, the acid and conjugate base pair for that step have equal concentrations.
Vocabulary
- Polyprotic acid
- An acid that can donate more than one proton per molecule in separate steps.
- Stepwise ionization
- The process in which a polyprotic acid loses its protons one at a time, with a different equilibrium constant for each step.
- Ka
- The acid dissociation constant that measures how strongly an acid donates a proton in a specific equilibrium step.
- Dominant species
- The acid or conjugate base form present at the highest concentration at a particular pH.
- Equivalence point
- The point in a titration where the moles of added base match the moles of acidic protons neutralized for a given stage.
Common Mistakes to Avoid
- Treating all protons as if they dissociate at the same time is wrong because each proton has its own Ka value and equilibrium step.
- Assuming Ka1, Ka2, and Ka3 are equal is wrong because removing each later proton is usually harder as the remaining species becomes more negatively charged.
- Using the total acid concentration as the concentration of every species is wrong because the acid is distributed among forms such as H3A, H2A-, HA2-, and A3- depending on pH.
- Missing multiple equivalence points in a titration is wrong when the acid has well-separated Ka values, because each removable proton can create a separate neutralization stage.
Practice Questions
- 1 Phosphoric acid has pKa1 = 2.15, pKa2 = 7.20, and pKa3 = 12.35. At what pH are [H2PO4-] and [HPO4 2-] equal?
- 2 A 25.0 mL sample of 0.100 M H2A is titrated with 0.100 M NaOH. What volume of NaOH is needed to reach the first equivalence point, and what volume is needed to reach the second equivalence point?
- 3 A triprotic acid H3A has pKa values of 2.0, 6.5, and 11.0. Explain which species is dominant at pH 4.0 and why it is between two ionization steps.