2.1 The correct indicator was added to 25.00 mL of a 0.1345 M solution of propanoic acid (CH3CH2COOH) (Ka = 1.3 x 10-5 ) in a flask, and this was titrated with a 0.1895 M solution of sodium hydroxide from a burette. 2.1.1 Write the net ionic balanced equation for the reaction taking place (include all phases) 2.1.2 Determine the final concentrations of the acid-base conjugate pair present after the titration, and show that the ratio of these species' concentrations is 1:1 at half the equivalence point. Hint: Firstly, determine the volume of base (in mL) that is required to reach the equivalence point for this titration. You may make use of a table to determine the amounts of species present. 2.1.3 Using the Henderson-Hasselbalch equation, determine the pH of the solution in the flask at this point.

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Question 2 – Propanoic acid
2.1 The correct indicator was added to 25.00 mL of a 0.1345 M solution of propanoic acid
(CH3CH2COOH) (Ka = 1.3 x 10-5 ) in a flask, and this was titrated with a 0.1895 M solution of
sodium hydroxide from a burette.
2.1.1 Write the net ionic balanced equation for the reaction taking place (include all phases)
2.1.2 Determine the final concentrations of the acid-base conjugate pair present after the
titration, and show that the ratio of these species' concentrations is 1:1 at half the
equivalence point.
Hint:
Firstly, determine the volume of base (in mL) that is required to reach the equivalence
point for this titration.
You may make use of a table to determine the amounts of species present.
2.1.3 Using the Henderson-Hasselbalch equation, determine the pH of the solution in the flask at
this point.
Transcribed Image Text:Question 2 – Propanoic acid 2.1 The correct indicator was added to 25.00 mL of a 0.1345 M solution of propanoic acid (CH3CH2COOH) (Ka = 1.3 x 10-5 ) in a flask, and this was titrated with a 0.1895 M solution of sodium hydroxide from a burette. 2.1.1 Write the net ionic balanced equation for the reaction taking place (include all phases) 2.1.2 Determine the final concentrations of the acid-base conjugate pair present after the titration, and show that the ratio of these species' concentrations is 1:1 at half the equivalence point. Hint: Firstly, determine the volume of base (in mL) that is required to reach the equivalence point for this titration. You may make use of a table to determine the amounts of species present. 2.1.3 Using the Henderson-Hasselbalch equation, determine the pH of the solution in the flask at this point.
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