Glossary
Acid dissociation constant (Ka)
An equilibrium constant that quantifies the extent to which a weak acid dissociates in water, with a larger Ka indicating a stronger acid.
Example:
A weak acid with a Ka of 1.8 x 10^-5 indicates it only partially dissociates in solution.
Acids
Substances that have a pH less than 7 when dissolved in water and are known as proton donors.
Example:
Lemon juice is an acid because it tastes sour and has a pH around 2-3.
Base dissociation constant (Kb)
An equilibrium constant that quantifies the extent to which a weak base dissociates or accepts protons in water, with a larger Kb indicating a stronger base.
Example:
Ammonia has a Kb value that reflects its ability to accept protons and form hydroxide ions in water.
Bases
Substances that have a pH greater than 7, often feel slippery, and are known as proton acceptors.
Example:
Baking soda dissolved in water forms a base that can neutralize acidic spills.
Buffers
Solutions that resist significant changes in pH upon the addition of small amounts of acid or base, typically composed of a weak acid and its conjugate base or a weak base and its conjugate acid.
Example:
Blood contains a buffer system that helps maintain its pH within a narrow, life-sustaining range.
Equivalence Point
The point in a titration where the moles of acid completely neutralize the moles of base, or vice versa, according to the stoichiometry of the reaction.
Example:
At the equivalence point of a strong acid-strong base titration, the pH is 7.0, indicating complete neutralization.
Free protons
Hydrogen ions (H+) that are readily available in a solution to react with other chemical species.
Example:
In a strong acid solution, there are many free protons available to react with a base.
Half-Equivalence Point
The point in a weak acid-strong base (or weak base-strong acid) titration where exactly half of the weak acid (or base) has been neutralized, and pH = pKa (or pOH = pKb).
Example:
At the half-equivalence point of a weak acid titration, the concentration of the weak acid equals the concentration of its conjugate base, making pH = pKa.
Henderson-Hasselbalch Equation
An equation, pH = pKa + log([A-]/[HA]), used to calculate the pH of a buffer solution or to determine the ratio of conjugate base to weak acid needed for a specific pH.
Example:
To prepare a buffer with a specific pH, a chemist would use the Henderson-Hasselbalch Equation to determine the necessary concentrations of the weak acid and its conjugate base.
Ion product of water (Kw)
The equilibrium constant for the autoionization of water, equal to [H+][OH-], which is 1.0 x 10^-14 at 25°C.
Example:
The ion product of water allows us to calculate [OH-] if [H+] is known, or vice versa, in any aqueous solution.
Neutral Solutions
Solutions that have a pH of 7.0, indicating an equal concentration of hydrogen (H+) and hydroxide (OH-) ions.
Example:
Pure distilled water is a neutral solution with a pH of 7.0 at 25°C.
Proton Acceptors
A term for bases, referring to their ability to accept hydrogen ions (H+) in a chemical reaction.
Example:
Ammonia (NH3) acts as a proton acceptor when it reacts with water to form NH4+ and OH-.
Proton Donors
A term for acids, referring to their ability to donate hydrogen ions (H+) in a chemical reaction.
Example:
In the reaction HCl → H+ + Cl-, hydrochloric acid acts as a proton donor.
Strong acids
Acids that completely dissociate in water, releasing all their hydrogen ions (H+) into the solution.
Example:
Hydrochloric acid (HCl) is a strong acid used in laboratories, fully ionizing to produce H+ and Cl- ions.
Strong bases
Bases that completely dissociate in water, releasing all their hydroxide ions (OH-) into the solution.
Example:
Sodium hydroxide (NaOH), a common drain cleaner, is a strong base that fully dissociates in water.
Titration Curves
Graphs that plot the pH of a solution against the volume of titrant added during a titration, illustrating the change in pH throughout the reaction.
Example:
Analyzing the shape of titration curves helps identify the equivalence point and determine the strength of the acid or base being titrated.
Titrations
A quantitative analytical method used to determine the concentration of an unknown solution (analyte) by reacting it with a solution of known concentration (titrant).
Example:
A chemist performs a titration to find the exact concentration of acetic acid in a sample of vinegar.
Weak acids
Acids that only partially dissociate in water, meaning only a fraction of their molecules release hydrogen ions (H+).
Example:
Acetic acid (CH3COOH), found in vinegar, is a weak acid that establishes an equilibrium between its dissociated and undissociated forms.
Weak bases
Bases that only partially dissociate in water, meaning only a fraction of their molecules accept hydrogen ions or release hydroxide ions.
Example:
Caffeine is a weak base that can accept protons, contributing to its bitter taste.
pH
A measure of the hydrogen ion concentration in a solution, calculated as the negative logarithm of [H+].
Example:
If a solution has a hydrogen ion concentration of 1.0 x 10^-4 M, its pH would be 4.0.
pH Formula
The mathematical equation pH = -log[H+], used to calculate the pH of a solution from its hydrogen ion concentration.
Example:
Using the pH Formula, if [H+] is 1.0 x 10^-7 M, the pH is 7.0.
pH Scale
A logarithmic scale ranging from 0 to 14 that measures the acidity or basicity of an aqueous solution based on its hydrogen ion concentration.
Example:
A solution with a pH of 1 is highly acidic, while a solution with a pH of 13 is highly basic, as indicated by the pH scale.
pOH
A measure of the hydroxide ion concentration in a solution, calculated as the negative logarithm of [OH-].
Example:
A solution with a hydroxide ion concentration of 1.0 x 10^-2 M would have a pOH of 2.0.
pOH Formula
The mathematical equation pOH = -log[OH-], used to calculate the pOH of a solution from its hydroxide ion concentration.
Example:
If a solution has an [OH-] of 1.0 x 10^-3 M, the pOH Formula yields a pOH of 3.0.