Acid Base Titration 2015

Acid-Base Titration
Background Information
A titration is a controlled addition of one substance into another substance. In an acid-base titration, the experimenter will add a base of known concentration to an acid of unknown concentration (or vice-versa). The goal of the titration is usually to use the substance of known concentration to determine the concentration of the other substance. In order to run a titration, the following materials are needed:
• A buret filled with the base (or acid) of known concentration
• A beaker or flask containing a measured volume of acid (or base) of unknown concentration
• Several drops of a chemical indicator, which will be added into the flask with the acid
The titration is performed by slowly dripping the basic solution from the burette into a beaker or flask containing a measured amount of the acidic solution and several drops of a chemical indicator. An indicator is a chemical that will respond to changes in its environment. In the case of an acid-base titration, the experimenter will most often use an indicator that will change color when the endpoint of the titration is reached. In an acid-base titration the experimenter is trying to determine the equivalence point of the reaction, which is the point when the amount of base added was exactly the correct amount to have the moles of base completely react with the moles of acid. If the experimenter chooses the correct indicator, the endpoint of the reaction will be as close as possible to the actual equivalence point. The indicator most commonly used for acid base titrations is phenolphthalein. Phenolphthalein is a colorless material in acidic and weakly basic solutions. Once the solution is titrated to a more basic solution the indicator turns pink.

Before titrating an unknown acid we will first standardize the NaOH solution. By this we mean we will accurately determine its concentration through a titration experiment. A known amount of a standard (we will use KHP) is titrated with the unknown NaOH solution. Using the stoichiometry of the acid-base reaction the molarity of the unknown NaOH solution can be calculated. KHP is potassium hydrogen phthalate, KNaC8H4O4 (molar mass = 204.22 g/mole). KHP is a monoprotic acid. The reaction between NaOH and KHP is as follows:

NaOH(aq) + KHC8H4O4(aq) → H2O(l) + KNaC8H4O4(aq)

Thus 1 mole of NaOH reacts exactly with 1 mole of KHP. We can determine the concentration of the NaOH as follows moles acid = moles base (mass KHP/204.22g/mol) = MBVB (where B denotes base) MB = (mass KHP/204.22)/VB (note VB must be in units of L)

After standardizing our NaOH solution we will use it to titrate a common household chemical: vinegar. Vinegar is a dilute aqueous solution of acetic acid (HC2H3O2). With the data generated we will calculate the molarity and the of the acetic acid in vinegar. The titration reaction has the following stoichiometry:

HC2H3O2(aq) + NaOH(aq) → H2O + NaC2H3O2(aq)

With a stoichiometry of 1 to 1 between acetic acid and sodium hydroxide, the equivalence point data allows the calculation of the molarity of the unknown solution with the following equation:

moles acid = moles base MA x VA = MB x VB (where A denotes acid and B denotes base)

The molarity of the base will have been previously determined from the standardization experiment. The volume of the base needed to reach the end point (VB) is read from the buret and the volume of the vinegar (VA) used will be measured at the start of the experiment. A rearrangement of the above equation will allow for the calculation of MA:

MA = (MB x VB) / VA
Procedure

Standardization of NaOH

a. Prepare a clean buret for use.
1. Rinse twice with deionized H2O
2. Rinse twice with a small amount (5 mL) of the NaOH solution.
b. Fill the buret with the NaOH solution. Record the volume.
c. Accurately weigh around 0.2000 – 0.4000 grams of KHP.
d. Place KHP in 125 mL Erlenmeyer flask.
e. Add 25 mL deionized H2O to the flask, swirl to dissolve KHP.
f. Add 2 – 3 drops of phenolphthalein solution.
g. Titrate KHP with the NaOH solution until you obtain a faint pink color that does not disappear. Be sure to add the NaOH solution slowly with constant swirling of the flask.
h. Record the volume of NaOH required to reach the end point.
i. Repeat steps (b) to (h) 2 more times.
Titration of Vinegar

a. Using a pipet, measure 10.00 mL of vinegar and place it into a 100.00 mL volumetric flask. Add water to the 100.00 mL mark on the flask. Mix well.
b. Pipet 10.00 mL of the diluted vinegar and place in a 125 mL Erlenmeyer flask.
c. Add 2 – 3 drops of phenolphthalein solution
d. Fill your buret with NaOH. Record the volume of NaOH.
e. Titrate the vinegar with the standardized NaOH until you obtain a faint pink color that does not disappear. Be sure to add the NaOH solution slowly with constant swirling of the flask
f. Record the volume of NaOH required to reach the end point.

Repeat steps (b) to (f) 2 more times.

Data Sheet

Standardization of NaOH Solution

Trial 1 Trial 2 Trial 3 Mass of KHP

Initial buret reading

Final buret reading

Volume NaOH used

Molarity of NaOH solution

Average NaOH molarity

Standard Deviation

Titration of Vinegar

Trial 1 Trial 2 Trial 3

Initial buret reading

Final buret reading

Volume NaOH used

Molarity of acetic acid in diluted vinegar:

Average molarity of diluted vinegar

Standard Deviation

Molarity of undiluted vinegar:

Post Lab Questions

Name

1. A 0.3423 g sample of KHP required 33.98 mL of a NaOH solution to reach the phenolphthalein end point. Determine the molarity of the NaOH solution.

2. 15.46 mL of the NaOH solution described in Question 1 was required to titrate 10.00 mL of acetic acid. Determine the concentration of the acetic acid.