Acid Titration Curve (Page 2 of 4) in pdf

II. Prelab (Practice in interpreting titration curves)

1. For each titration curve, find (be precise in a & b!):

a. The equivalence point volume and pH.

Acid 1

b. The half equivalence point volume and pH.

c. The K

of the acid being titrated.

2. Which acid is the weaker acid? How do you know?

Acid 2

3. Which sample had more moles of HA in it initially?

Give reason!

4. Do the prelab calculation found in part IIIA below.

5. Calculate the (expected) initial concentration of your acid

solution (see Part IIIB below to get the initial volume of your

acid solution).

(mL)

Volume of titrant (base) added

III. Experimental Information and Procedures

A. Background and preliminary calculation. You will be given access to one of the following four acids:

(i) Mandelic Acid (C

CH(OH)COOH; 152.16 g/mol); (ii) Glycolic Acid (CH

(OH)COOH; 76.05 g/mol);

(an acidic salt; 204.23 g/mol); (iv) NaH

O (an acidic salt; 138.0 g/mol)

(iii) KHC

•H

You will be preparing a solution containing your acid, and then you will titrate that solution with a

standardized solution of approximately 0.1 M NaOH (“standardized” means that its concentration is

known with great accuracy and precision; record the value from the NaOH(aq) reagent bottle).

Prelab Calculation: Given the molar mass of your acid (just pick one), calculate the number of grams of

the acid needed to just react with 30.0 mL of 0.100 M NaOH.

Now, given what you just calculated, if you were to start with that number of grams of acid in a titration,

at what “volume of titrant added” would you expect to reach the equivalence point if the titrant used were

0.100 M NaOH? (Note this is NOT a trick question, I just want to make sure you understand what the

meaning of “equivalence point” is in a titration!) I hope that you realize that the answer to this question is

“30.0 mL”! Please ask me if you don’t understand this.

B. Preparation of the acid solution.

1. TARE a clean and dry 100-mL beaker on the balance (i.e., put the beaker on the pan and press the

TARE button; the reading should go to 0.000 g), and then weigh out the number of grams of your acid

that you calculated in part A into the beaker (to within 0.01 g of what you calculated; do not waste

time trying to get it to match more precisely than that).

2. Add 40.0 mL of deionized water using a graduated cylinder, and swirl to dissolve all of the solid.

3. Add a magnetic stir bar, and possibly a few drops of an acid-base indicator (optional; ask me).

C. Titration of the acid solution.

1. Basic ideas and Setup.

a. You will be adding NaOH solution (titrant) in small aliquots (from a buret) to your acid solution and

measuring and recording the pH after each addition of titrant. Once you are finished with the

titration (as indicated below), you will plot “pH vs. volume of titrant added (in mL)”, so make sure

you have plenty of room on a separate sheet of paper to record your values of “volume of titrant

added” and corresponding pH (make a table, with headings; THIS TABLE WILL BE HANDED IN).

You should also write down the concentration of the NaOH solution from the bottle at this point.

Obtain approximately 75 mL of the titrant (NaOH) solution in a clean and dry 150 mL Erlenmeyer

flask (Note: It is always a good ideas to swirl the reagent bottle of a standardized solution before

removing any solution). Place a watch glass on this flask when not pouring from it.

b. Buret Prep. Before the titration, prepare the buret by doing the following: 1) Make sure that the

stopcock is firmly seated in the end of the buret! There are two types of tips, and if you have the

wrong one inserted, it could fall out when you start turning the stopcock! 2) Rinse the buret with 2

or 3 portions of ~3 mL of the titrant (NaOH) solution (as I’ll demonstrate). 3) Attach the buret to a

buret clamp on a ring stand, and fill the buret to above the 0.00 mL line with titrant. 4) Bring the

Acid Titration Curve Page 2

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