Acids and Bases

1. Acids and Bases Titrations AP Chemistry

2. Neutralization Reactions and Titrations • Neutralization Reactions • Strong acid + Strong Base  Salt + Water • HCl + NaOH  NaCl + H2O (neutral pH = 7) • Neutral solution results only if the acid and base are mixed in the mole ratios specified in the balanced equation.

3. Titration • A lab technique used to determine the unknown concentration of an acid or base using a neutralization reaction.

4. Titration Set-up • Titrant- an acid/base with a known concentration • Analyte- a solution being analyzed • Burette- glassware used to deliver the titrant

5. For example, if you had an acid of unknown concentration, you would add base (with a known concentration) to it until all the acid had been neutralized. • At the very point that all the acid was gone, the amount of base you added would exactly equal the amount of acid you started with. • Moles of Acid = Moles of Base

6. The equation for titration calculations is • At neutralization, Moles of Acid= Moles of Base MaVa = MbVb Ma = Acid Molarity Mb = Base Molarity Va = Acid volume Vb = Base Volume • Volumes can be any units as long as they are the same on both sides!

7. Example Titration Problem • If 5.00mL of 1.00M NaOH was used to titrate10.0mL of HNO3, what is the concentration of HNO3? MaVa = MbVb Ma=? Va=10.0mL Mb=1.00M Vb= 5.00mL Ma (10.0mL) = (1.00M)(5.00mL) Ma = (1.00M)(5.00mL) (10.00mL) Ma =0.500M

8. Titration Vocabulary • Equivalence Point – • The point in a titration where the moles of acid = moles of base • Endpoint – Point at which the indicator changes color. • The endpoint and the equivalence points are not always the same. (but they should be!!!) • An indicator is used to determine when the acid has been neutralized in a titration. • Without an indicator it would be impossible to know when the titration should stop, unless you use a pH meter.

9. Titration of strong acid with a strong base • Such as NaOH and HCl • Equivalence point occurs at pH = 7 • Should use an indicator that changes at a pH of 7, such as bromothymol blue

10. Strong Acid – Strong Base Titrations • You can calculate the volume of base needed to reach the equivalence point using the formula: V·M = V·M. • There are three situations in which you determine pH. • Initial strong acid concentration (this is simply the –log[H+] which is based on the [Acid].) • Equivalence point (or endpoint) when moles of OH- = moles of H+. The pH is 7 (due to the auto-ionization of water.) • Before and after the endpoint (calculate excess moles of H+ or OH-, divide by the total volume, and calculate the pH based on this value.)

11. Titration of a weak acid with a strong base • Such as CH3COOH and NaOH • Equivalence point occurs at a pH >7 • Should use an indicator that changes at a pH above 7, such as phenolphthalein. • Notice that starting pH is higher and • the Equivalence Point pH is higher.

12. Why is the Equivalence point at a pH >7? The conjugate base of a weak acid is a strong base. • CH3COOH = Weak acid • CH3COO- = Strong Base • Once all of the acid has been converted to CH3COO-, it starts to take the H+ from H2O in solution, creating more OH- ions thus making the pH >7

13. Weak Acid – Strong Base Titrations • When a weak acid (such as HC2H3O2) is neutralized by a strong base (such as NaOH), the graph varies in two ways: • the equivalence point is not at pH = 7 and • a buffer region exists as you approach the endpoint. • You can still calculate the volume of base needed to reach the equivalence point using the formula: V·M = V·M. Weak acids require the same amount of base for neutralization as strong acids because they dissociate as they are neutralized

14. Weak Acid – Strong Base Titrations Buffer regions Halfway to equivalence point pH = pKa

15. There are five situations in which you need to be able to calculate the pH. • 1. Initial weak acid concentration (this is an ICE box calculation.) The shortcut can be used here. • 2. Equivalence point (endpoint) is when all of the weak acid has been neutralized and turned into the conjugate base (C2H3O2- in this case.) This is a hydrolysis problem. Calculate the [C2H3O2-] and then do an ICE box problem knowing that Kb = . Calculate the [OH-], the pOH, and then the pH. • 3. Halfway to the equivalence point (as in a half-titration) the pH = pKa. This is because at this point, there is a perfect buffer as the [HA] = [A-]. At this point, you can determine the Ka of an unknown weak acid… very useful.

16. 4. Before and after the half-way point, the pH can be calculated using the Henderson-Hasselbach equation (or an ICE box, if you want.) Use stoichiometry to determine the [HA] and [A-]. pH = pKa + log • 5. Finally, after the equivalence point, the pH depends on the excess strong base that has been added. As in the strong acid-strong base titration, calculate excess moles of OH-, divide by the total volume, and calculate the pOH and then pH based on this value. The effect on the pH by the A- is negligible compared to the excess OH-.

17. Weak Base – Strong Acid Titrations • When a sample of a weak base is titrated with a strong acid, the curve resembles an inverted Weak Acid – Strong Base titration curve. • Note that the pH at the equivalence point is less than 7. • An indicator such as phenolphthalein that changes at pH of 9 would change when only 6 mL of acid had been added even though the equivalence point is reached at around 11 mL. The acid-base indicator must be chosen with a Ka near to the [H+] of the equivalence point; that is the pKa of the indicator must match the pH of the equivalence point.

18. Weak Diprotic Acid – Strong Base Titrations • When a weak diprotic acid (examples: H2C2O4 or H2CO3) is titrated, there are two equivalence points. • The curve is not as distinct because of the various proton donors and proton acceptors in solution.