Water's Amphoteric Nature and Autoionization

Study Notes

The autoionization of water has the following equilibrium:
- H2O (l) + H2O (l) ↔ H3O+ (aq) + OH– (aq)

The ion product constant for water (Kw) is defined as:
- Kw = [H3O+] [OH–] = 1.00 × 10−14 at 25 °C

All aqueous solutions contain both H3O+ and OH– ions.
The product of the concentrations of H3O+ and OH– is always 1.00 × 10−14 at room temperature.

Neutral solutions have an equal concentration of H3O+ and OH–.
- [H3O+] = [OH–] = 1.00 × 10−7 M

Acidic solutions have a higher concentration of H3O+ than OH–.
- [H3O+] > 1.00 × 10−7 M; [OH–] < 1.00 × 10−7 M

Basic solutions have a higher concentration of OH– than H3O+.
- [H3O+] < 1.00 × 10−7 M; [OH–] > 1.00 × 10−7 M

pOH measures the hydroxide ion concentration:
- pOH = −log[OH−]
- [OH−] = 10−pOH
The sum of pH and pOH is always 14.0
- pH + pOH = 14.0

Strong acids fully dissociate in water, meaning they donate all their protons (H+) to water molecules, forming hydronium ions (H3O+).
Example: In a 0.10 M HNO3 solution, HNO3 completely dissociates into 0.10 M H3O+ and 0.10 M NO3-.
To calculate pH, use the formula: pH = -log[H3O+]. In the HNO3 example, pH = -log(0.10) = 1.
Percent dissociation is 100% for strong acids.

Weak acids only partially dissociate in water, establishing an equilibrium between the undissociated acid and its conjugate base.
To find the Ka value of a weak acid, use the equation: Ka = ([H3O+][A-])/[HA], where [HA] is the initial concentration of the weak acid, [H3O+] is the equilibrium concentration of hydronium ions, and [A-] is the equilibrium concentration of the conjugate base.
You can calculate the pH of a weak acid solution using the Ka value and the initial concentration of the acid.
Percent dissociation for a weak acid is calculated as: (Concentration of dissociated acid)/(Initial concentration of acid) * 100%.

In a mixture of weak acids, the stronger acid (with a larger Ka value) will dissociate more than the weaker acid.
The dissociation of the stronger acid will inhibit the dissociation of the weaker acid, as the equilibrium of the weaker acid will shift to the left (towards the undissociated form) due to the increased concentration of H3O+ ions.
To estimate the pH of a mixture of weak acids, you need to consider the contribution of both acids to the total H3O+ concentration. This may require an iterative calculation, as the dissociation of each acid will affect the equilibrium of the other.