Acid-Base Titrations - AP Chemistry

Graph of pH versus volume of titrant added. Visual representation showing pH changes during titration process.

To determine the concentration of an unknown solution. Uses known titrant to find unknown concentration via stoichiometry.

$M = \frac{\text{moles of solute}}{\text{liters of solution}}$. Defines concentration as moles of solute per liter of solution.

Where moles of acid equal moles of base. Stoichiometric neutralization point on titration curve.

pH = 7. Equal amounts of strong acid and base neutralize to neutral pH.

Resists changes in pH upon addition of acid or base. Contains conjugate acid-base pair to maintain pH stability.

Increases the pH. Dilution decreases $[H^+]$, raising pH toward neutral.

8.2 to 10.0. Indicator's transition range for color change.

Weak acid-strong base. Weak acid-strong base gives basic equivalence point (pH > 7).

50 mL. Equal molarity and volume means 1:1 mole ratio requires equal volume.

The solution added from the burette. Standard solution delivered from burette during analysis.

A solution of known concentration used in titration. Standard solution added from burette to determine analyte concentration.

0.0025 moles. Convert volume to L (0.025 L), multiply by molarity (0.1 M).

Not always; depends on the strength of acid and base. Only true for strong acid-strong base; weak acids/bases differ.

$[\text{H}^+] = 10^{-\text{pH}}$. Inverse relationship: pH is negative log of $[H^+]$

Phenolphthalein. Changes color at pH 8.2-10, ideal for strong acid-base endpoints.

The point at which the indicator changes color. Visual signal used to determine when to stop adding titrant.

To deliver a precise volume of titrant. Graduated tube allows controlled addition and volume measurement.

The point where the indicator changes color. Observable color change indicating approximate equivalence point.

Higher $K_a$ indicates a stronger acid. Larger $K_a$ means greater ionization and stronger acid.

pH = 11. pOH = $-\log(10^{-3}) = 3$, so pH = $14 - 3 = 11$.

$\text{pH} + \text{pOH} = 14$. Fundamental relationship for aqueous solutions at 25°C.

Slight decrease in pH. Buffer components consume added acid, minimizing pH change.

$[\text{OH}^{-}] = 10^{-\text{pOH}}$. Inverse relationship: pOH is negative log of $[OH^-]$.

To signal the endpoint of the titration. Color change marks approximate equivalence point.

pH = 6. pH is defined as $-\log[H^+] = -\log(10^{-6}) = 6$.

Neutralization. Acid and base react to form salt and water.

pOH = 2. Strong base: pOH = $-\log(0.01) = 2$.

To measure and transfer a specific volume of solution. Precision instrument for accurate volume transfer.