1. Fill burette with titrant. 2. Measure analyte into flask. 3. Add indicator. 4. Slowly add titrant, swirling, until indicator changes color.

React a known volume of the unknown solution (analyte) with a solution of known concentration (titrant) until the reaction reaches the equivalence point, then use stoichiometry to calculate the unknown concentration.

1. Identify the pairs in the reaction. 2. The compound with the extra hydrogen is the acid; the other is the base.

1. Identify knowns. 2. Use the equation $M_aV_a = M_bV_b$. 3. Plug in and solve, remembering to convert volumes to liters and account for mole ratios.

Solution with a known concentration, usually in the burette.

Solution with an unknown concentration, usually in the Erlenmeyer flask.

The point where moles of titrant = moles of analyte; reaction is complete.

The point where the indicator changes color, ideally close to the equivalence point.

Graph showing pH change of the analyte as titrant is added.

A proton (H⁺) donor.

A proton (H⁺) acceptor.

The species remaining after an acid has donated a proton.

The species formed when a base accepts a proton.

A substance that can act as both an acid and a base.

Equivalence Point: Moles of acid equals moles of base (theoretical). Endpoint: Indicator changes color (experimental).

Strong Acid/Strong Base: No buffer region, sharp equivalence point. Weak Acid/Strong Base: Buffer region present, equivalence point at pH > 7.

Acid: Proton donor. Conjugate Base: What remains after the acid donates a proton. Differ by one proton.