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 M solution of sodium chloride g NaCl Mix Thoroughly!!

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

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

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 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

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 = M HCl Concentration of HCl is now known to 3 Sig. Figs. instead of just 1 !! 5) Calculate average Molarity and standard deviation.