Diprotic Acids and Bases: Videos & Practice Problems

Diprotic acids, such as carbonic acid, can donate two hydrogen ions (H⁺), leading to two distinct dissociation constants, K_a1 and K_a2. K_a1 is always greater than K_a2, with K_a1 representing the loss of the first hydrogen ion and K_a2 the second. The intermediate form of the acid, after donating the first hydrogen ion, is an anion (e.g., HCO₃⁻), and after donating the second, it becomes a fully realized base (e.g., CO₃²⁻). The relationship between K_a and the corresponding K_b is such that K_a1 is paired with K_b2 and K_a2 with K_b1, with both products equaling the ion product constant for water (K_w). Equilibrium expressions for these reactions are determined by the ratio of the concentrations of the products over reactants. Understanding these concepts is crucial for analyzing the behavior of diprotic acids in various chemical reactions and their equilibrium states.

Ka values of Diprotic Acids

Diprotic acids are characterized by their ability to donate two protons (H⁺) in solution, which is represented by the general formula H₂A. This property allows them to have two distinct acid dissociation constants, denoted as K_a1 and K_a2. It is important to note that K_a1 is always greater than K_a2, indicating that the first proton is more readily donated than the second.

The first dissociation can be expressed as:

H₂A ⇌ H⁺ + HA^-   (K_a1)

In this reaction, K_a1 represents the equilibrium constant for the dissociation of the first acidic proton. The second dissociation follows as:

HA^- ⇌ H⁺ + A^2-   (K_a2)

Here, K_a2 corresponds to the equilibrium constant for the donation of the second proton. Understanding these dissociation processes is crucial as they relate to the concepts of acid strength and the equilibrium expressions for diprotic acids. The relationship between K_a and K_b (the base dissociation constant) also plays a significant role in determining the behavior of these acids in various chemical equilibria.

3 forms of Diprotic Acids

In the study of diprotic acids, understanding the dissociation process is crucial. A diprotic acid, such as carbonic acid (H₂CO₃), can donate two protons (H⁺) in a stepwise manner. Initially, in its fully protonated form, H₂A represents the acid with both acidic protons attached. The first dissociation step, characterized by the acid dissociation constant K_a1, involves the release of the first H⁺ ion, resulting in the intermediate form, HA^-. This intermediate still retains one acidic proton and is crucial for understanding the acid's behavior in solution.

As the process continues, the second dissociation occurs, represented by K_a2, where the remaining H⁺ ion is released, transforming the species into A^2-, which is now a fully deprotonated base. This transition highlights the importance of recognizing the acid-base equilibrium and the role of K_a values in determining the strength of the acid at each stage.

Conversely, when considering the reverse process, the base A^2- can accept protons, leading to the formation of HA^- and H₂A. The constants K_b1 and K_b2 describe these base dissociation reactions. The relationships between these constants are significant: K_a1 is related to K_b2, and K_a2 is related to K_b1, with both pairs equating to the ion product of water (K_w).

For example, in the case of carbonic acid, the first dissociation can be expressed as:

Ka₁ = \frac{[HCO₃^-][H₃O^+2CO₃]}.

Following this, the second dissociation is represented as:

Ka₂ = \frac{[CO₃^2-][H₃O^+3-]}.

These equilibrium expressions illustrate the dynamic nature of diprotic acids and the importance of understanding both the acidic and basic forms. As such, careful attention must be paid to the specific K_a and K_b values associated with each step of the dissociation process, as they dictate the behavior of the acid in various chemical environments.

Diprotic Acids and Bases Example

Carbonic acid (H₂CO₃) is classified as a weak diprotic acid, meaning it can donate two protons (H⁺) in solution. The dissociation of carbonic acid occurs in two steps, characterized by the acid dissociation constants K_a1 and K_a2. When carbonic acid loses both of its protons, it forms the carbonate ion (CO₃^2-), which represents the basic form of the species.

In this basic form, the carbonate ion can react with water, where water acts as an acid by donating a proton. This reaction produces the bicarbonate ion (HCO₃^-) and hydroxide ions (OH^-). The equilibrium constant for this reaction is known as the base dissociation constant (K_b1).

To relate K_b1 to the acid dissociation constants, we use the relationship between K_a2 and K_b1, which is defined by the equation:

$K_b=\frac{K}{w}\frac{K}{a}$

Where K_w is the ion product of water, equal to 1.0 × 10^-14 at 25°C. Given that K_a2 is 5.6 × 10^-11, we can calculate K_b1 as follows:

$K_b=\frac{K}{w}\frac{K}{a}$

Calculating this gives:

$K_b=\mathrm{1.8\; \backslash times\; 10^\{-4\}}$

Thus, the base dissociation constant K_b1 for the carbonate ion is 1.8 × 10^-4.

Determine the equilibrium expression for the K_a2 of hydrosulfuric acid, H₂S?

• A. K_a2 = [HS⁻][H₃O⁺]/[H₂S][H₂O]
• B. K_a2 = [S²⁻][H₃O⁺]/[HS⁻][H₂O]
• C. K_a2 = [HS⁻][H₃O⁺]/[H₂S]
• D. K_a2 = [S²⁻][H₃O⁺]/[HS⁻]

Answer: (D)

Here’s what students ask on this topic:

What is a diprotic acid?

A diprotic acid is a type of acid that has the capacity to donate two protons or hydrogen ions (H⁺) per molecule in an aqueous solution. This donation of protons occurs in a stepwise manner, with the first proton typically being released more easily than the second. The dissociation of a diprotic acid happens in two distinct stages, each characterized by its own dissociation constant (K_a). The first dissociation constant (K_a1) corresponds to the release of the first proton, resulting in the formation of a monoprotic anion. The second dissociation constant (K_a2) is for the release of the second proton, leading to the formation of a diprotic anion. Examples of diprotic acids include sulfuric acid (H₂SO₄), carbonic acid (H₂CO₃), and oxalic acid (C₂H₂O₄). Each of these acids undergoes two separate ionization reactions in water, reflecting their diprotic nature.

What is an example of a diprotic acid?

A diprotic acid is an acid that can donate two protons or hydrogen ions (H⁺) per molecule in an aqueous solution. An example of a diprotic acid is sulfuric acid (H₂SO₄). In the first stage of ionization, sulfuric acid donates a single proton to form the hydrogen sulfate ion (HSO₄^-), which makes it a monoprotic acid at this point. However, the hydrogen sulfate ion can further donate a second proton to form the sulfate ion (SO₄^2-), which demonstrates the diprotic nature of sulfuric acid. This two-step dissociation process is a characteristic of diprotic acids, where each proton is released in a separate reaction, often with different dissociation constants. The first dissociation constant (K_a1) is typically larger than the second (K_a2), indicating that the first proton is more easily lost than the second.