Glossary
'p' Notation
A mathematical shorthand used to express very small numbers, typically concentrations, as more manageable positive values by taking the negative logarithm of the value.
Example:
Instead of writing [H+] = 1.0 x 10^-7 M, we use pH = 7.0, which is much simpler to work with.
Acid Strength
A measure of how readily an acid donates a proton (H+) in solution. Stronger acids dissociate more completely, while weaker acids dissociate only partially.
Example:
Hydrochloric acid (HCl) exhibits high acid strength because it fully dissociates in water, unlike acetic acid (CH3COOH), which only partially dissociates.
Acid-Base Indicator
A weak acid or base that changes color over a specific pH range, used to visually determine the equivalence point of an acid-base titration.
Example:
Phenolphthalein is a common acid-base indicator that turns pink in basic solutions and remains colorless in acidic solutions.
Buffer
A solution that resists significant changes in pH upon the addition of small amounts of acid or base. It typically consists of a weak acid and its conjugate base, or a weak base and its conjugate acid.
Example:
Blood acts as a natural buffer system, maintaining a stable pH essential for biological processes.
Equivalence Point
The point in a titration where the moles of titrant added are stoichiometrically equal to the moles of the analyte initially present.
Example:
In a strong acid-strong base titration, the equivalence point occurs at pH 7, where the moles of H+ equal the moles of OH-.
Half-equivalence point
The point in a titration of a weak acid (or base) where exactly half of the initial weak acid (or base) has been neutralized by the titrant. At this point, [weak acid] = [conjugate base], and pH = pKa.
Example:
During the titration of a weak acid, the half-equivalence point is reached when the volume of added base is half of the volume needed to reach the equivalence point.
Henderson-Hasselbalch Equation
An equation that relates the pH of a buffer solution to the pKa of the weak acid and the ratio of the concentrations of the conjugate base and weak acid.
Example:
Using the Henderson-Hasselbalch Equation, a student can calculate the pH of a buffer made from acetic acid and sodium acetate.
pH
A measure of the hydrogen ion concentration ([H+]) in a solution, indicating its acidity or alkalinity. It is calculated as -log[H+].
Example:
A solution with a pH of 2 is highly acidic, while a solution with a pH of 12 is strongly basic.
pKa
The negative logarithm of the acid dissociation constant (Ka), which quantifies the strength of a weak acid. A lower pKa indicates a stronger acid.
Example:
Acetic acid has a pKa of about 4.76, making it a weak acid, whereas hydrochloric acid is a strong acid and doesn't typically have a pKa value listed in this context.
pKb
The negative logarithm of the base dissociation constant (Kb), which quantifies the strength of a weak base. A lower pKb indicates a stronger base.
Example:
Ammonia has a pKb of about 4.75, indicating it is a weak base.
pOH
A measure of the hydroxide ion concentration ([OH-]) in a solution, indicating its basicity. It is calculated as -log[OH-].
Example:
If a solution has a pOH of 3, it means it's very basic, and its pH would be 11 at 25°C.