Go back in the chapter and examine the values of the equilibrium constants for the three dissociations of the triprotic phosphoric acid. Give a reason for why each successive dissociation yields a weaker acid. (Hint: Acids must give up a proton, and opposite charges attract.)

An equilibrium constant, represented as Ka for acid dissociation, plays a critical role in understanding the behavior of acids in water. It serves as a quantifiable measure of the extent to which an acid dissociates in a solution. This dissociation is a reversible process, and at equilibrium, the rates of the forward and reverse reactions are equal. The equilibrium constant is calculated by taking the concentration of the products (dissociated ions) and dividing it by the concentration of the reactants (undissociated acid), raised to the power of their respective stoichiometric coefficients.

For a general acid dissociation reaction:HA ⇌ H+ + A-, the acid dissociation constant is:
Ka = [H+][A-]/[HA]

The larger the Ka value, the stronger the acid, indicating a higher degree of dissociation into protons (H+) and anions (A-). Consequently, a higher Ka value implies a lower pH of the solution, as more H+ ions are present. Students should remember that typically, each successive acid dissociation step will have a smaller Ka value, reflecting a reduced tendency to release additional protons and thus a weaker acid strength.

Triprotic phosphoric acid, H3PO4, is an acid capable of donating three protons (H+) in a sequential manner. This means phosphoric acid can undergo three distinct dissociation steps, with each step corresponding to the release of one proton. The molecule is constructed of one phosphorus atom surrounded by four oxygen atoms - three of which are hydroxyl groups (with hydrogen atoms) and one being a double-bonded oxygen atom.

During the dissociation process, phosphoric acid behaves in the following manner:

• First dissociation: H3PO4 becomes H2PO4-
• Second dissociation: H2PO4- becomes HPO42-
• Third dissociation: HPO42- becomes PO43-

After each dissociation step, the negative charge of the remaining ion increases, leading to a decrease in the acid strength because positive and negative charges attract each other. This interplay of charges is fundamental in predicting the reactivity and acid strength of the remaining ions after each proton is lost.

Dissociation reactions are processes where molecules or ionic compounds separate into smaller particles, often ions. In the context of acids, a dissociation reaction involves the release of hydrogen ions (protons, H+) from the acid molecule. Each dissociation step for an acid has its unique position of equilibrium and accompanying equilibrium constant.

For triprotic acids like phosphoric acid, the dissociation can be represented by a sequence of reactions:

• First dissociation: H3PO4 ⇌ H+ + H2PO4-
• Second dissociation: H2PO4- ⇌ H+ + HPO42-
• Third dissociation: HPO42- ⇌ H+ + PO43-

As acid dissociates, the nature of the remaining ion changes – essentially becoming less acidic with each successive proton released. This trend is a critical concept students need to grasp to understand acid strength across multiple dissociations.

Chemical equilibrium refers to the state of a reversible chemical reaction where the rate of the forward reaction equals the rate of the reverse reaction. As a result, the concentrations of reactants and products remain constant over time, even though both reactions continue to occur. Equilibrium does not mean that the reactants and products are present in equal amounts, but that their ratios do not change.

In the case of acids, particularly triprotic acids like phosphoric acid, achieving chemical equilibrium in each dissociation step is essential to understand their behavior in solution. The position of equilibrium is what determines the acid strength and is indicated by the Ka value. It's vital to note that while at equilibrium the concentration of species remains constant, the system can still respond to changes in conditions according to Le Chatelier's Principle, which could lead to shifts in the position of equilibrium.