Glossary
Buffer region (of a titration curve)
The portion of a titration curve where the pH changes very slowly upon the addition of titrant, indicating the presence of a significant amount of both the weak acid and its conjugate base.
Example:
On a titration curve of a weak acid with a strong base, the flat part before the steep rise is the buffer region, where the solution resists pH changes.
Buffers
Solutions that resist significant changes in pH upon the addition of small amounts of acid or base. They consist of a weak acid and its conjugate base, or a weak base and its conjugate acid.
Example:
Our blood contains natural buffers that help maintain a stable pH, preventing drastic changes that could be harmful to bodily functions.
Conjugate base ([A-])
The species formed when a weak acid donates a proton (H+). In a buffer, it works with the weak acid to neutralize added acid.
Example:
In an acetic acid buffer, the acetate ion (CH3COO-) acts as the conjugate base, ready to react with any added H+ ions.
Equivalence point
The point in a titration where the moles of titrant added are stoichiometrically equal to the moles of analyte initially present.
Example:
When titrating a strong acid with a strong base, the equivalence point occurs at pH 7, indicating complete neutralization.
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, and the concentration of the weak acid equals the concentration of its conjugate base. At this point, pH = pKa.
Example:
At the half-equivalence point during the titration of acetic acid, the pH of the solution will be exactly equal to acetic acid's pKa.
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:
When calculating the pH of a blood sample, chemists often use the Henderson-Hasselbalch equation to account for the bicarbonate buffer system.
ICE table
An acronym for Initial, Change, Equilibrium; a tabular method used to organize and calculate equilibrium concentrations of reactants and products in a reversible reaction.
Example:
To find the equilibrium concentrations of ions in a weak acid dissociation, an ICE table helps systematically track the changes from initial amounts to equilibrium.
Ka
The acid dissociation constant, an equilibrium constant that quantifies the extent to which a weak acid dissociates in solution. A larger Ka indicates a stronger acid.
Example:
The Ka value for hydrofluoric acid (HF) is 6.8 × 10⁻⁴, which is relatively large for a weak acid, indicating it dissociates more than many other weak acids.
Kb
The base dissociation constant, an equilibrium constant that quantifies the extent to which a weak base dissociates in solution. It is related to Ka by Kw = Ka * Kb.
Example:
If you know the Kb of ammonia, you can calculate the pH of an ammonia solution, as it tells you how much hydroxide ion is produced.
Net Ionic Equation
A chemical equation that shows only the species that participate directly in the reaction, excluding spectator ions.
Example:
For the reaction of hydrochloric acid and sodium hydroxide, the net ionic equation is H+(aq) + OH-(aq) → H2O(l), showing only the ions that form water.
Stoichiometry
The quantitative relationship between reactants and products in a chemical reaction, often expressed in moles or molar ratios.
Example:
Using stoichiometry, a chemist can predict how much product will be formed from a given amount of reactants, or how much reactant is needed for a desired amount of product.
Titration
A quantitative chemical analysis method used to determine the concentration of an identified analyte by reacting it with a precisely known concentration of a reagent.
Example:
A chemist performs a titration to determine the exact concentration of an unknown acid solution by slowly adding a base of known concentration until the reaction is complete.
Weak acid ([HA])
An acid that only partially dissociates in water, meaning it does not donate all of its protons. It is a key component of a buffer system.
Example:
Citric acid, found in oranges, is a weak acid that contributes to their tart taste and can be part of a buffer system.
pH
A measure of the acidity or alkalinity of a solution, calculated as the negative logarithm (base 10) of the hydrogen ion concentration ([H+]).
Example:
A lemon has a low pH of around 2, indicating its high acidity, while baking soda dissolved in water has a high pH, making it basic.
pKa
The negative logarithm (base 10) of the acid dissociation constant (Ka), which indicates the strength of a weak acid. A lower pKa means a stronger acid.
Example:
Acetic acid has a pKa of 4.74, which is why it's a weak acid and a common component in buffer solutions.