Polyprotic Acids — Solve polyprotic acid problems (intro) and identify dominant species - General Chemistry 2

SECTION 1

Historical Context & Motivation

The chemistry of acids capable of donating more than one proton has been central to understanding biological, geological, and industrial systems for well over a century. Early investigations into polyprotic acids — acids possessing two or more ionizable hydrogen atoms — revealed that these species do not surrender all their protons simultaneously. Instead, each proton departs in a distinct, thermodynamically characterized step, a finding that reshaped our understanding of acid–base equilibria and solution speciation. The development of accurate dissociation models for polyprotic acids proved essential for everything from analyzing carbonate buffering in blood plasma to predicting mineral solubility in natural waters.

1884 • Arrhenius Electrolytic Theory

Svante Arrhenius proposed that acids dissociate in water to produce H⁺ ions. His framework laid the groundwork for quantifying acid strength but treated multi-proton acids as simple extensions of monoprotic behavior.

1909 • Sørensen Introduces pH

Søren Sørensen defined the pH scale, providing a practical logarithmic measure of hydrogen-ion concentration. This tool became indispensable for tracking the stepwise proton release from polyprotic acids.

1923 • Brønsted–Lowry Definition

Johannes Brønsted and Thomas Lowry independently redefined acids as proton donors and bases as proton acceptors, providing the conjugate-pair framework essential for analyzing each successive dissociation of a polyprotic acid.

1940s • Quantitative Ka Measurements

Advances in potentiometric titration allowed precise determination of successive dissociation constants (Ka₁, Ka₂, Ka₃) for common polyprotic acids such as phosphoric acid and citric acid, confirming that each Ka drops by several orders of magnitude.

1960s–Present • Alpha-Fraction and Speciation Diagrams

With computational chemistry and graphing tools, scientists routinely generate alpha (α) diagrams that display the fraction of each protonation state as a function of pH, enabling rapid identification of the dominant species at any given acidity.

The central question this lesson addresses is straightforward yet powerful: given a polyprotic acid at a known concentration, how do we calculate the pH of the solution, and how do we determine which protonation state dominates at a particular pH? Answering these questions requires understanding stepwise dissociation equilibria, the relative magnitudes of successive dissociation constants, and the concept of dominant species in solution.

SECTION 2

Core Principles & Definitions

Before solving numerical problems, it is essential to internalize several foundational ideas that govern polyprotic acid behavior. A polyprotic acid is any Brønsted–Lowry acid that can donate more than one proton per molecule. Common examples include sulfuric acid (H₂SO₄, diprotic), carbonic acid (H₂CO₃, diprotic), and phosphoric acid (H₃PO₄, triprotic). Each successive proton loss constitutes its own equilibrium with a unique acid-dissociation constant, and the interplay among these constants dictates which species predominates at any pH.

Stepwise Dissociation

Polyprotic acids release protons one at a time, not all at once. Each step has its own equilibrium expression and its own K_a value. We write separate balanced equations for each dissociation.

Decreasing Ka Values

Each successive K_a is smaller than the one before it — typically by a factor of 10³ to 10⁵. Removing a proton from an increasingly negative ion requires more energy, so each step is thermodynamically less favorable.

First Dissociation Dominates pH

Because K_a1 ≫ K_a2, the [H⁺] from the first step overwhelms subsequent contributions. We can usually approximate the pH using only the first dissociation.

Dominant Species Depends on pH

At low pH the fully protonated form dominates; at high pH the fully deprotonated form dominates. Intermediate species dominate in pH ranges between successive pK_a values. The crossover points occur at pH = pK_a.

The Simplifying Assumption

When K_a1/K_a2 > 10³, we treat each dissociation independently. The [H⁺] produced in step 1 is assumed unaffected by step 2, and [A²⁻] ≈ K_a2 for diprotic acids.

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KEY TAKEAWAY

Think of a polyprotic acid as a multi-stage rocket: each stage fires independently and with decreasing thrust. The first stage (K_a1) delivers the vast majority of propulsion (H⁺ ions), the second stage (K_a2) adds a comparatively tiny additional push, and a potential third stage (K_a3) barely nudges the trajectory. Because the stages are so unequal, we can almost always analyze the dominant stage alone and obtain an accurate result.

SECTION 3

Visual Explanation — Stepwise Dissociation of Phosphoric Acid

Phosphoric acid (H₃PO₄) serves as the canonical triprotic acid for illustrating stepwise dissociation. The diagram below depicts each successive proton loss, the corresponding conjugate base produced, and the relative magnitude of each dissociation constant. Notice how each K_a decreases by roughly five orders of magnitude, visually reinforcing why the first dissociation overwhelmingly determines solution pH.

The top row shows the four protonation states of phosphoric acid connected by arrows representing each proton-loss step. The bottom row summarizes the K_a and pK_a for each step, illustrating that each constant drops by approximately five orders of magnitude.

The diagram makes the critical pattern unmistakable: the electrostatic penalty of removing a proton from an already-negatively-charged ion increases dramatically with each step. Removing H⁺ from neutral H₃PO₄ is relatively easy (K_a1 ≈ 10⁻³), but removing H⁺ from the doubly-charged HPO₄²⁻ is extraordinarily difficult (K_a3 ≈ 10⁻¹³). This enormous disparity in successive K_a values is what enables the simplifying assumption that we can treat each dissociation step independently.

SECTION 4

Mathematical Framework

We now formalize the equilibrium expressions for a general diprotic acid H₂A and develop the mathematical strategy for computing pH and identifying species concentrations. The approach extends naturally to triprotic acids.

FIRST DISSOCIATION

H₂A ⇌ H⁺ + HA⁻ Ka₁ = [H⁺][HA⁻] / [H₂A]

K_a1 is the equilibrium constant for the loss of the first proton from the fully protonated species H₂A.

SECOND DISSOCIATION

HA⁻ ⇌ H⁺ + A²⁻ Ka₂ = [H⁺][A²⁻] / [HA⁻]

K_a2 is always smaller than K_a1. The conjugate base HA⁻ from step 1 acts as the acid in step 2.

pH APPROXIMATION (STEP 1 ONLY)

Ka₁ = x² / (C₀ − x) where x = [H⁺] ≈ [HA⁻]

C₀ is the initial concentration of H₂A. Because K_a1 ≫ K_a2, we treat the problem as a monoprotic acid equilibrium. If x < 5% of C₀, further simplify to x ≈ √(K_a1 × C₀).

CONCENTRATION OF A²⁻

[A²⁻] ≈ Ka₂ (for diprotic acids when Ka₁ ≫ Ka₂)

This remarkably simple result arises because [H⁺] ≈ [HA⁻] from step 1, so K_a2 = [H⁺][A²⁻]/[HA⁻] simplifies to [A²⁻] ≈ K_a2 when [H⁺] ≈ [HA⁻]. This confirms that the second dissociation barely alters [H⁺].

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When to Use the Quadratic

If K_a1 is relatively large compared to C₀ (i.e., x > 5% of C₀), the small-x approximation fails and you must solve the full quadratic: x² + K_a1x − K_a1C₀ = 0. This situation commonly arises for moderately strong first dissociations like those of oxalic acid or phosphoric acid at low concentration.

SECTION 5

Identifying Dominant Species — Alpha Fractions & pH Regions

One of the most practically important skills with polyprotic acids is the ability to identify the dominant species at a given pH. The dominant species is the protonation state present at the highest concentration. For a diprotic acid H₂A with pK_a1 and pK_a2, the rule is elegantly simple: when pH < pK_a1, the fully protonated form H₂A dominates; when pK_a1 < pH < pK_a2, the intermediate species HA⁻ dominates; and when pH > pK_a2, the fully deprotonated form A²⁻ dominates. Transitions occur at pH = pK_a, where two adjacent species are present in equal concentrations.

Alpha-fraction diagram for the carbonic acid system. The cyan curve represents the fraction of total dissolved carbonate present as H₂CO₃, the violet curve as HCO₃⁻, and the pink curve as CO₃²⁻. The yellow dashed lines mark the pK_a values where adjacent species are present in equal concentrations.

The alpha-fraction diagram above uses the carbonic acid/bicarbonate system — the critical buffer in human blood — as an example. At physiological pH (≈ 7.4), which falls between pK_a1 = 6.35 and pK_a2 = 10.33, the bicarbonate ion (HCO₃⁻) is overwhelmingly the dominant species. This is precisely why the bicarbonate buffer system is so effective: at blood pH, there is ample HCO₃⁻ available to neutralize both added acid and added base.

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Quick Dominant-Species Rule

Compare the given pH to each pK_a in order. Every time pH exceeds a pK_a, the acid has "lost" that proton. Count how many pK_a values the pH exceeds — that tells you how many protons have been removed, and hence which species dominates.

SECTION 6

Worked Example — pH of a Diprotic Acid Solution

Let us work through a complete problem: find the pH, [H⁺], [HA⁻], and [A²⁻] for a 0.10 M solution of ascorbic acid (vitamin C, H₂C₆H₆O₆), given K_a1 = 8.0 × 10⁻⁵ and K_a2 = 1.6 × 10⁻¹². We abbreviate the acid as H₂A for clarity.

pH and Species Concentrations for 0.10 M Ascorbic Acid

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Step 1 — Confirm the simplifying assumption applies

Check whether K_a1/K_a2 > 10³. Here, (8.0 × 10⁻⁵)/(1.6 × 10⁻¹²) = 5.0 × 10⁷, which is vastly greater than 10³. We can therefore treat the two dissociations independently.

K_a1/K_a2 = 5.0 × 10⁷ ≫ 10³ ✓
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Step 2 — Solve the first dissociation as a monoprotic acid

Set up the ICE table for H₂A ⇌ H⁺ + HA⁻. Let x = [H⁺] = [HA⁻]. Then K_a1 = x²/(0.10 − x). Try the small-x approximation: x ≈ √(8.0 × 10⁻⁵ × 0.10) = √(8.0 × 10⁻⁶) = 2.83 × 10⁻³ M.

x = 2.83 × 10⁻³ M
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Step 3 — Validate the 5% approximation

Check: (2.83 × 10⁻³ / 0.10) × 100% = 2.83%. Since 2.83% < 5%, the small-x approximation is valid.

2.83% < 5% — approximation valid ✓
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Step 4 — Calculate pH

pH = −log[H⁺] = −log(2.83 × 10⁻³) = 2.55.

pH = 2.55
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Step 5 — Determine [HA⁻] and [A²⁻]

From step 2, [HA⁻] = x = 2.83 × 10⁻³ M. For [A²⁻], we use the result that [A²⁻] ≈ K_a2 when the simplifying assumption holds: [A²⁻] ≈ 1.6 × 10⁻¹² M. This confirms that the second dissociation contributes a negligible amount of H⁺ to the solution.

[HA⁻] = 2.83 × 10⁻³ M; [A²⁻] = 1.6 × 10⁻¹² M
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Step 6 — Identify the dominant species

Compare: [H₂A] = 0.10 − 0.00283 ≈ 0.097 M ≫ [HA⁻] = 2.83 × 10⁻³ M ≫ [A²⁻] = 1.6 × 10⁻¹² M. The dominant form of the acid in solution is H₂A, the fully protonated species. This makes sense because pH = 2.55 is well below pK_a1 = 4.10.

Dominant species: H₂A (fully protonated ascorbic acid)

SECTION 7

Comparing Common Polyprotic Acids

Different polyprotic acids vary significantly in the number of ionizable protons, the magnitude and spacing of their K_a values, and their biological or industrial importance. The table below summarizes several key examples and highlights an important exception: sulfuric acid, whose first dissociation is strong (K_a1 is very large) while the second is weak.

Acid	Formula	Ka₁	Ka₂	Ka₃	Notes
Sulfuric acid	H₂SO₄	Strong (≫1)	1.2 × 10⁻²	—	Step 1 is complete; treat step 2 as weak acid
Carbonic acid	H₂CO₃	4.3 × 10⁻⁷	4.7 × 10⁻¹¹	—	Blood buffer system; HCO₃⁻ dominates at pH 7.4
Phosphoric acid	H₃PO₄	7.5 × 10⁻³	6.2 × 10⁻⁸	4.8 × 10⁻¹³	Triprotic; used in fertilizers and soft drinks
Oxalic acid	H₂C₂O₄	5.9 × 10⁻²	6.4 × 10⁻⁵	—	Large Ka₁; quadratic often required
Ascorbic acid	H₂C₆H₆O₆	8.0 × 10⁻⁵	1.6 × 10⁻¹²	—	Vitamin C; classic simplification example

Dissociation constants for selected polyprotic acids at 25 °C

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KEY TAKEAWAY

Sulfuric acid is the important exception to the standard polyprotic workflow. Because its first dissociation is essentially complete, you should treat it as a strong acid that produces 0.10 M H⁺ and 0.10 M HSO₄⁻ initially, and then solve for the additional [H⁺] from the weak second dissociation of HSO₄⁻ (K_a2 = 1.2 × 10⁻²). Think of it as a hybrid: the first dissociation is "automatic" like a strong acid, while the second dissociation requires an ICE table.

SECTION 8

Connection to Advanced Acid–Base Theory

The introductory treatment presented here — using the simplifying assumption that K_a1 ≫ K_a2 to decouple successive equilibria — works remarkably well for most polyprotic acids encountered in general chemistry. However, more advanced situations demand a fuller treatment. In analytical chemistry and biochemistry, you will encounter systems where the simplification breaks down or where the pH is not simply determined by one dissociation.

Feature	Introductory Approach	Advanced Approach
Equilibria considered	Each step independently	All steps simultaneously (systematic treatment of equilibrium)
pH of amphoteric species (e.g., NaHA)	Not typically covered	pH ≈ ½(pKa₁ + pKa₂)
Alpha-fraction expressions	Qualitative (from diagrams)	Quantitative: α₀, α₁, α₂ expressed as functions of [H⁺], Ka₁, Ka₂
Titration curves	Recognize buffering regions and equivalence points	Full calculation of pH at every point using mass/charge balance
Applicability	Ka₁/Ka₂ > 10³ (most polyprotic acids)	Any Ka ratio; required for amino acids, EDTA, etc.

Introductory vs. advanced treatment of polyprotic acid systems

The concept of amphoteric species deserves special mention. When a salt like NaHCO₃ or Na₂HPO₄ dissolves, it produces an intermediate species that can act as either an acid or a base. The pH of such solutions is approximately pH ≈ ½(pK_a1 + pK_a2), a remarkably elegant result that is independent of concentration. You will derive this formula in analytical chemistry courses, but understanding dominant species and stepwise dissociation from this lesson provides the conceptual foundation you need.

SECTION 9

Practice Problems

Question 1: CONCEPTUAL

Explain why K_a2 is always smaller than K_a1 for any polyprotic acid. What fundamental electrostatic argument underlies this trend?

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The second proton must be removed from a species that already carries a negative charge (HA⁻), which means the departing H⁺ is electrostatically attracted to the conjugate base more strongly than it was in the neutral molecule. This increased Coulombic attraction raises the energy required for dissociation, making K_a2 smaller. The same logic applies to each subsequent dissociation: removing a proton from a −2 ion is even harder than from a −1 ion. This electrostatic argument explains the universal trend K_a1 > K_a2 > K_a3.

Question 2: BASIC CALCULATION

Calculate the pH of a 0.050 M solution of carbonic acid (H₂CO₃). K_a1 = 4.3 × 10⁻⁷ and K_a2 = 4.7 × 10⁻¹¹.

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Since K_a1/K_a2 ≈ 9.1 × 10³ > 10³, we treat only the first dissociation. Using x ≈ √(K_a1 × C₀) = √(4.3 × 10⁻⁷ × 0.050) = √(2.15 × 10⁻⁸) = 1.47 × 10⁻⁴ M. The 5% check: (1.47 × 10⁻⁴/0.050) × 100% = 0.29%, so the approximation is valid. pH = −log(1.47 × 10⁻⁴) = 3.83. The dominant species of the carbonate system at this pH is H₂CO₃ because pH = 3.83 < pK_a1 = 6.37.

Question 3: INTERMEDIATE

For a 0.20 M solution of oxalic acid (H₂C₂O₄), K_a1 = 5.9 × 10⁻² and K_a2 = 6.4 × 10⁻⁵. Determine [H⁺] and explain whether you need the quadratic formula.

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First, K_a1/K_a2 = 922, which is close to but slightly below 10³. We proceed with step 1 only but should note the approximation is borderline. The small-x approximation gives x ≈ √(5.9 × 10⁻² × 0.20) = √(1.18 × 10⁻²) = 0.109 M. Check: 0.109/0.20 = 54% — this far exceeds 5%, so we must use the quadratic. x² + 0.059x − 0.0118 = 0. Using the quadratic formula: x = [−0.059 + √(0.059² + 4 × 0.0118)] / 2 = [−0.059 + √(0.0507)] / 2 = [−0.059 + 0.225] / 2 = 0.083 M. Thus [H⁺] = 0.083 M and pH = 1.08.

Question 4: APPLIED

A water sample from a limestone aquifer has a pH of 8.3 and contains dissolved carbonate species. Using the pK_a values for carbonic acid (pK_a1 = 6.35, pK_a2 = 10.33), identify the dominant carbonate species and explain what role it plays in the natural buffering of this water.

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At pH = 8.3, we compare with the two pK_a values: 6.35 < 8.3 < 10.33. Since the pH lies between pK_a1 and pK_a2, the dominant species is HCO₃⁻ (bicarbonate). Bicarbonate acts as an amphoteric buffer: it can accept protons (from acid rain or metabolic acid inputs) to form H₂CO₃ and release CO₂, or it can donate protons to bases to form CO₃²⁻. This dual capability makes the bicarbonate system a critical natural buffer that maintains the pH of natural waters within a range habitable for aquatic organisms.

Question 5: CRITICAL THINKING

Consider a hypothetical diprotic acid H₂X with K_a1 = 1.0 × 10⁻³ and K_a2 = 5.0 × 10⁻⁴. Would the standard simplifying assumption (treating each dissociation independently) be valid? Discuss what complications arise and how you would approach calculating the pH of a 0.10 M solution.

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The ratio K_a1/K_a2 = 2.0, which is far less than 10³. This means the second dissociation contributes significantly to [H⁺], and the approximation [H⁺] ≈ [HX⁻] is no longer valid because a substantial fraction of HX⁻ further dissociates to X²⁻. One would need to solve the system of two coupled equilibrium expressions simultaneously, along with the mass balance ([H₂X] + [HX⁻] + [X²⁻] = 0.10) and charge balance ([H⁺] = [HX⁻] + 2[X²⁻]). This is the systematic treatment of equilibrium, which typically requires solving a cubic or quartic equation or using numerical/iterative methods. This scenario mirrors real systems like certain amino acids where the two pKₐ values are close together.

SUMMARY

Lesson Summary

Summary

Polyprotic acids donate protons in stepwise dissociations, each characterized by its own acid-dissociation constant (Ka). Because each successive K_a is typically 10³ to 10⁵ times smaller than the previous one, the first dissociation dominates the pH of the solution. This allows us to treat the first step as a monoprotic acid equilibrium — setting up an ICE table, solving for x = [H⁺], validating the 5% approximation, and using the quadratic formula when necessary.

To identify the dominant species at a given pH, compare pH to each pKa value in sequence: the number of pK_a thresholds exceeded tells you how many protons have been lost. Alpha-fraction diagrams provide a powerful visual tool for this analysis. Remember that [A²⁻] ≈ Ka₂ for diprotic acids under the simplifying assumption, and that sulfuric acid is the notable exception requiring a hybrid strong-acid/weak-acid approach. These concepts form the foundation for understanding buffer systems, polyprotic titrations, and the speciation chemistry of biological and environmental systems.