Acid and Base Titration Lecture Notes

Comprehensive vocabulary flashcards covering acid-base titration principles, indicators, standard solutions, complex systems, and analytical methods like Kjeldahl analysis based on lecture notes.

Titration

An analytical technique that links laboratory practice with chemical quantification by comparing the moles of analyte with the moles of titrant.

Titrant

A standard solution with an accurately known concentration placed in the burette, such as $0.1000\,M\,NaOH$.

Analyte

The unknown solution placed in the conical flask to be analyzed, such as a vinegar sample.

Standard Solution

A solution with an accurately known concentration, prepared directly from a primary standard or obtained by standardization.

Equivalence Point

The theoretical point where the amount of acid and base reacted exactly according to the balanced equation stoichiometry.

End Point

The experimental point where the chosen indicator changes color, such as phenolphthalein turning pale pink.

Concordant Titres

Repeated titration results that are obtained within a close range, commonly within $\pm 0.10\,mL$.

Parallax Error

A reading bias caused by failing to record burette readings at eye level.

Indicator

A weak acid or base ($HIn$) whose color depends on pH and the ratio $[In^-]/[HIn]$.

Methyl Orange

An indicator with a pH transition range of $3.1\text{--}4.4$ that changes color from red to yellow.

Bromothymol Blue

An indicator with a pH transition range of $6.0\text{--}7.6$ that changes color from yellow to blue.

Phenolphthalein

An indicator with a pH transition range of $8.2\text{--}10.0$ that changes color from colorless to pink.

Primary Standard

A substance characterized by high purity, stability in air, being non-hygroscopic, having a high molar mass, and reacting completely with known stoichiometry.

Potassium Hydrogen Phthalate (KHP)

A primary standard commonly used for the standardization of sodium hydroxide ($NaOH$).

Standardization

The process of determining the true molarity of a solution, such as $NaOH$, by reacting it against a primary standard to account for impurities or air absorption.

Buffer Region

A region in a weak acid/base titration where the solution resists pH change due to the presence of the weak species and its conjugate.

Half-Equivalence Point

The point in a weak acid titration where $[HA] = [A^-]$, resulting in $pH = pK_a$.

Henderson-Hasselbalch Equation

The formula used to calculate pH in a buffer region: $pH = pK_a + \log\left(\frac{[A^-]}{[HA]}\right)$.

Polyprotic Acid

An acid containing more than one ionizable hydrogen atom, such as $H_2CO_3$ or $H_3PO_4$.

Zwitterion

The form of an amino acid ($^+H_3N\text{--}CHR\text{--}COO^-$) at an intermediate pH where it has no net charge.

Isoelectric Point (pI)

The pH at which an amino acid has no net charge, calculated as $pI = \frac{pK_{a1} + pK_{a2}}{2}$ for neutral amino acids.

Kjeldahl Analysis

A method used to determine the nitrogen content in organic samples to estimate protein content.

Protein Conversion Factor (F)

A factor used to convert nitrogen percentage to protein percentage; a typical value is $6.25$, assuming $16\%\text{ nitrogen}$.

Arrhenius Acid/Base

A definition where an acid produces $H^+/H_3O^+$ in water and a base produces $OH^-$ in water.

Br%nsted–Lowry Acid/Base

A definition where an acid donates a proton and a base accepts a proton.

Lewis Acid/Base

A definition where an acid accepts an electron pair and a base donates an electron pair.

Potentiometric Titration

An automated titration method using a pH electrode to detect the endpoint, offering higher precision and suitability for colored or turbid samples.