Triprotic Acids and Bases: Videos & Practice Problems

Triprotic acids, such as phosphoric acid, have the ability to donate three protons (H⁺ ions), resulting in three distinct acid dissociation constants (K_a values). The K_a values decrease with each proton lost, as it becomes increasingly difficult to lose subsequent protons. The first dissociation constant (K_a1) is associated with the loss of the first proton, K_a2 with the second, and K_a3 with the third. The basic form of a triprotic acid, after losing all three protons, can accept protons in reverse, leading to three corresponding base dissociation constants (K_b values). These K_b values are inversely related to the K_a values, with K_a1 related to K_b3, K_a2 to K_b2, and K_a3 to K_b1, each pair multiplying to equal the ion product constant for water (K_W). Understanding these relationships and the equilibrium expressions for each dissociation step is crucial in analyzing the behavior of triprotic acids in various chemical contexts.

Triprotic Acid Ka Values

Triprotic acids are characterized by the general formula H₃A, indicating that they can donate three acidic protons (H⁺ ions). Each of these protons corresponds to a distinct acid dissociation constant, represented as K_a1, K_a2, and K_a3. The magnitudes of these constants follow a specific trend: K_a1 is always greater than K_a2, which in turn is greater than K_a3. This pattern arises because the process of losing protons becomes increasingly difficult with each successive proton lost.

Specifically, K_a1 pertains to the dissociation of the first proton, K_a2 relates to the second proton, and K_a3 corresponds to the final proton. Understanding which K_a value is relevant is crucial when analyzing the behavior of triprotic acids, as it directly influences the acid's strength and reactivity in various chemical contexts. This knowledge is essential for predicting the acid's behavior in solution and its interactions with other chemical species.

Triprotic Acid Equilibrium

Triprotic acids, also known as polyprotic acids, are characterized by their ability to donate three protons (H⁺) in a stepwise manner. This process results in four distinct forms of the acid as it loses its acidic protons. The initial form, H₃A, retains all three acidic protons, representing the fully protonated state. Upon donating the first H⁺, it transforms into H₂A^-, which is associated with the first dissociation constant, K_a1. The second dissociation leads to the formation of HA^2-, linked to K_a2, while the final dissociation results in A^3-, corresponding to K_a3. This last form is the basic state, having lost all acidic protons.

The relationship between the acid dissociation constants (K_a) and the base dissociation constants (K_b) is crucial in understanding the behavior of triprotic acids. For instance, K_a1 is related to K_b3, K_a2 to K_b2, and K_a3 to K_b1. The product of each pair of K_a and K_b values equals the ion product of water, K_w, which is defined as K_a × K_b = K_w.

To illustrate these concepts, consider phosphoric acid (H₃PO₄), a common triprotic acid. The first dissociation can be represented as:

H₃PO₄ + H₂O ⇌ H₂PO₄^- + H₃O⁺ (K_a1)

In this reaction, phosphoric acid donates its first proton to water, forming dihydrogen phosphate and hydronium ions. The equilibrium expression for this reaction is:

K_a1 = \frac{[H₂PO₄^-][H₃O^+3PO₄]}

The second dissociation involves the intermediate form H₂PO₄^- reacting with water:

H₂PO₄^- + H₂O ⇌ HPO₄^2- + H₃O⁺ (K_a2)

And its equilibrium expression is:

K_a2 = \frac{[HPO₄^2-][H₃O^+2PO₄-]}

Finally, the third dissociation can be expressed as:

HPO₄^2- + H₂O ⇌ PO₄^3- + H₃O⁺ (K_a3)

With the equilibrium expression:

K_a3 = \frac{[PO₄^3-][H₃O^+42-]}

Understanding these dissociation steps and their corresponding equilibrium expressions is essential for analyzing the behavior of triprotic acids in various chemical contexts. The complexity of triprotic acids arises from their multiple dissociation steps, each with its own K_a value, making it important to track which K_a or K_b is relevant in a given scenario.

Triprotic Acid Dissociation Example

The dissociation equation for the second ionization of pyrophosphoric acid (H₃P₂O₇) can be represented as follows:When the first hydrogen ion (H⁺) is lost, the species present is H₃P₂O₇⁻. This species then donates a second hydrogen ion to water, which acts as a base. The reaction can be written as:\[\text{H}_3\text{P}_2\text{O}_7^- (aq) + \text{H}_2\text{O} (l) \rightleftharpoons \text{H}_2\text{P}_2\text{O}_7^{2-} (aq) + \text{H}_3\text{O}^+ (aq)\]In this equation, H₃P₂O₇⁻ donates a proton (H⁺) to water, resulting in the formation of H₂P₂O₇²⁻ and hydronium ions (H₃O⁺). This reaction is associated with the acid dissociation constant \( K_a2 \), which quantifies the strength of the acid in its second dissociation step.

Determine the equilibrium expression for the K_a3 value of citric acid, H₃C₆H₅O₇?

• A. K_a3 = [H₂C₆H₅O₇⁻][H₃O⁺]/[H₃C₆H₅O₇]
• B. K_a3 = [HC₆H₅O₇²⁻][H₃O⁺]/[H₂C₆H₅O₇⁻]
• C. K_a3 = [C₆H₅O₇³⁻][H₃O⁺]/[HC₆H₅O₇²⁻]
• D. K_a3 = [C₆H₅O₇³⁻][H₃O⁺]/[HC₆H₅O₇²⁻][H₂O]

Answer: (C)

Here’s what students ask on this topic:

What is triprotic acid?

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Triprotic acid is a hypothetical term and does not refer to a specific chemical compound in standard scientific nomenclature. However, the term "triprotic" suggests an acid with three proton-donating groups, meaning it can donate three hydrogen ions (<sup>H+</sup>) when dissolved in water. Acids with multiple ionizable hydrogen atoms are known as polyprotic acids, and a triprotic acid would be a type of polyprotic acid.

In a broader sense, if an acid can donate three protons, it undergoes dissociation in stages, with each hydrogen ion being released sequentially. With each proton donation, the acid becomes a different species: first, the intact acid (<sub>H3</sub>A), then the dihydrogen ion (<sub>H2</sub>A<sup>-</sup>), followed by the hydrogen ion (HA<sup>2-</sup>), and finally the triply-deprotonated ion (A<sup>3-</sup>). Each of these species would exist in equilibrium in a solution, with the position of the equilibrium depending on the pH of the solution and the acid's dissociation constants (Ka values).

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