1. Steps in preparing a solution of known molar concentration: 250 mL 1) Decide how much (volume) solution you want to make and what concentration (Molarity) you want it to be. 2) If starting with a solid solute, calculate how much (grams) to add to the appropriate volumetric flask before filling to the line with solvent. Example: I want to prepare 250 mL of a 0.100 M solution of sodium chloride. 1.46 g NaCl Mix Thoroughly!!

2. Arrhenius Acid – yields H 3 O + ions (hydronium ions) in aqueous solution Terminology Associated with Acid – Base Titrations Arrhenius Base – yields OH - ions (hydroxide ions) in aqueous solution Example: HCl (aq) + H 2 O (l)  H 3 O + (aq) + Cl - (aq) Example: NaOH (aq)  Na + (aq) + OH - (aq) Titrant – solution being delivered from the buret (typically of known concentration) Analyte – solution being titrated (typically of unknown concentration) Endpoint – the point in the titration where you stop titrating (mols acid = mols base) Indicator – typically a substance which changes color at the endpoint

3. When mols of titrant = mols of analyte, solution turns pink and you stop titrating.

4. Steps for standardizing a solution of approximate concentration in order to determine a much more accurate concentration. 1) Choose an appropriate primary standard compound. 2) Calculate how much to weigh out in order to require 30-40 mL of solution to reach the end point. Example: Sodium carbonate is often used as a primary standard for standardizing solutions of hydrochloric acid. The balanced equation for the reaction is shown below. = 0.19 g Na 2 CO 3

5. 3) Perform 3 replicate trials using accurately weighed (3 Sig Figs) amounts of Na 2 CO 3. Steps for standardizing a solution of approximate concentration in order to determine a much more accurate concentration. (cont.) Example: 4) Calculate more accurate Molarity = 0.101 M HCl Concentration of HCl is now known to 3 Sig. Figs. instead of just 1 !! 5) Calculate average Molarity and standard deviation.