 # Acids and Bases Titrations AP Chemistry. Neutralization Reactions and Titrations Neutralization Reactions Strong acid + Strong Base  Salt + Water HCl.

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Acids and Bases Titrations AP Chemistry

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

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

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

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

The equation for titration calculations is At neutralization, Moles of Acid= Moles of Base M a V a = M b V b 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!

Example Titration Problem If 5.00mL of 1.00M NaOH was used to titrate10.0mL of HNO 3, what is the concentration of HNO 3 ? M a V a = M b V b M a =? V a =10.0mLM b =1.00M V b = 5.00mL M a (10.0mL) = (1.00M)(5.00mL) M a = (1.00M)(5.00mL) (10.00mL) M a =0.500M

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.

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

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.)

Titration of a weak acid with a strong base Such as CH 3 COOH 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.

Why is the Equivalence point at a pH >7? The conjugate base of a weak acid is a strong base. CH 3 COOH = Weak acid CH 3 COO - = Strong Base Once all of the acid has been converted to CH 3 COO -, it starts to take the H+ from H 2 O in solution, creating more OH- ions thus making the pH >7

Weak Acid – Strong Base Titrations When a weak acid (such as HC 2 H 3 O 2 ) 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

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

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 (C 2 H 3 O 2 - in this case.) This is a hydrolysis problem. Calculate the [C 2 H 3 O 2 -] 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.

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-.

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.

Weak Diprotic Acid – Strong Base Titrations When a weak diprotic acid (examples: H 2 C 2 O 4 or H 2 CO 3 ) is titrated, there are two equivalence points. The curve is not as distinct because of the various proton donors and proton acceptors in solution.

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