DETERMINATION OF PERCENT KHP AND ACID EQUIVALENT WEIGHT
DETERMINATION OF PERCENT KHP AND ACID EQUIVALENT WEIGHT
Titration is a process of mixing measured volumes of reacting solutions in such a manner that one can determine when chemically equivalent amounts of reactants are mixed. One of the purposes of the titration process is to determine the concentration of a solute in a solution. Additionally, the titration process will be used in the analyses of soluble solid unknown acids.
The equivalence point of a titration is the point at which stoichiometric amounts of reactants have been mixed. A method must be used to show when the equivalence point has been reached. In acid-base titrations, phenolphthalein is often used as an indicator. Phenolphthalein is an organic molecule that is colorless in acidic solution and pink to red in the presence of base. If the indicator is placed in an acidic solution it will be colorless. As base is added to this solution, a pink color develops as the neutralization (equivalence) point is passed.
In this experiment each student will work alone and:
a. prepare an approximately 0.1-M NaOH solution.
b. standardize (precisely determine the molarity, ±0.0001 M) the NaOH using the pure solid monoprotic acid standard, potassium hydrogen phthalate (KHP).
c. determine the percent by mass of KHP in an impure sample.
d. determine the mass of a solid acid unknown that neutralizes one mole of hydroxide ion.
e. prepare a formal write-up of the experiment.
potassium hydrogen phthalate, (KHP)
Note: You will need to prepare your own data tables for this experiment. These will be turned in as part of the formal lab report. These tables must be prepared before you come to lab to begin data collection. You need to make sure that all measurements made have a place in the data tables (i.e., initial buret reading, final buret reading, etc.)
PROCEDURE
A. Preparation of an NaOH solution
1. Thoroughly clean your large screw cap bottle. Also clean its cap.
2. Review calculations for dilution and calculate the amount of 6-M NaOH needed to prepare 500 mL of approximately 0.1-M NaOH.
3. Using the 6-M NaOH in the hood, measure out the calculated amount of NaOH (use a graduated cylinder) and place it in the clean screw cap bottle.
4. Fill the bottle to the 500 mL line with distilled water (this is approximately 500 mL).
5. Cap the bottle and mix well by inversion (at least 20 inversions).
6. Put your name and/or locker number on the bottle of NaOH.
7. For this experiment you will need to create your own data tables. Read through the procedure (Parts B, C, and D) and create data tables for recording your data in the days to follow. You must have places in your data tables for all data taken in the procedure.
B. Standardization of NaOH solution
Note: The NaOH you have prepared is approximately 0.1-M. You must determine the precise molarity of your NaOH to at least 4 significant figures (keeping only the number of significant figures allowed).
1. Obtain and clean the buret assigned to your lab locker according to the signs posted in the lab. Rinse it 3 times with 2-3 mL of your NaOH solution prior to filling it. If you wish to use a beaker or funnel to help fill your burette you must clean them and then rinse them with your NaOH solution prior to their use. Clean a 250 mL Erlenmeyer flask (it does not have to be dry).
2. Fill the burette with your NaOH solution, rinse solution through the burette tip to eliminate air bubbles, and note the initial burette reading to two decimal places.
3. Obtain a capped vial of pure potassium hydrogen phthalate (KHP), standard. Label this vial and keep it capped when it is not in use.
4. Calculate the mass of pure KHP (molar mass = 204.23 g mol-1) that will require about 20 – 25 mL of approximately 0.1-M sodium hydroxide solution for complete reaction. Remember that KHP is monoprotic.
5. Do one trial. Take the vial of pure KHP, your clean flask, and the data sheet to the analytical balance room and measure KHP into the flask. Use the approximate mass (+/- 0.05 g) calculated in step 5 as a guide. Record the precise mass of KHP dispensed into the flask to the nearest 0.0001 g.
6. Dissolve the KHP in the flask in about 50 mL of distilled water.
7. Add 2 to 3 drops of phenolphthalein and titrate the flask to a consistent very, very faint pink end point. Record the final burette reading and calculate the total volume of NaOH used for the titration. The contents of the flask can now be discarded. (Save your KHP for additional trials.)
Note: If your titration volume was at least 10.00 mL (4 significant figures) this titration can be included in your calculations. However, a larger titration volume (closer to 25 mL) will give better precision. On the other hand, an unnecessarily large titration volume (more than 25 mL) is time consuming. The volume of NaOH solution required is directly proportional to the mass of KHP titrated. If the volume for your first your titration was not between 20 and 25 mL, adjust the mass used for the rest of your trials.
8. Clean three 250 mL Erlenmeyer flasks (they do not have to be dry) and label them #1, #2, and #3.
9. Take the vial of KHP, your 3 flasks, and your data sheet to the analytical balance room and measure KHP into each of the 3 flasks using the first trial as a guide. Record the exact mass of KHP in each flask to the nearest 0.0001 g.
Note: You should refill the buret for each titration. The NaOH solution remaining in your buret at the end of each lab session should be saved in a clean dry beaker and used for rinsing the buret at the next lab session. Never put unused solution back into your stock bottle. You risk contaminating your NaOH solution.
10. Follow steps 7 and 8 for each of your flasks.
11. Use the volume of NaOH solution and the mass of KHP in each flask to calculate the molarity of your NaOH. (You will need to average at least 3 values.)
12. Determine and record the average molarity of your NaOH. This solution will be used to determine the values for your unknowns. Take good care of it!!
13. Using at least three molarity values calculate your percent relative average deviation (see Appendix A at the end of this lab manual). Note: Percent relative average deviation is a measure of precision and at least 3 trials are required for the calculation to be meaningful. If your average deviation is less than 2%, it means that the data you have collected shows good precision and you have completed enough trials. If it is greater than 2%, then additional trials are needed.
C. Determination of percent KHP in an impure sample
1. Clean, rinse, and fill the buret with your NaOH as you have done for previous titrations.
2. Obtain a clean, dry capped shell vial containing an impure KHP unknown. Record your unknown’s number and label the vial. Keep this vial in your locker until your graded lab report has been returned to you.
Note: The shell vial of unknown contains enough sample for at least six trials. No additional unknown will be provided! Should an unknown be spilt, a different unknown will be obtained and you will start that unknown’s analysis from the beginning.
3. Do one trial titration with the unknown using about twice as much mass as was used for pure KHP. Record the precise mass of unknown (±0.0001 g) and volume of NaOH used (±0.01 mL). (Titration procedure is exactly the same as that used previously.)
4. You will need to do at least two more trials. If the total volume of NaOH used in your first titration was less than 20 mL use a little more unknown for your subsequent titrations. If your titration volume was greater than 25 mL use a little less unknown. (The mass of impure KHP and the volume of NaOH solution used in the titration are directly proportional).
5. Calculate the percent by mass of KHP in your impure sample for each trial.
6. Determine the percent relative average deviation using the calculated mass percents from all your trials. Do additional trials if your deviation is greater than 2%.
7. Report the average percent by mass of KHP for your unknown.
D. Determination of the mass of an unknown acid required to neutralize one mole of Hydroxide ion.
1. Obtain a clean, dry capped shell vial containing an impure KHP unknown. Record your unknown’s number and label the vial. Keep this vial in your locker until your graded lab report has been returned to you.
Note: the shell vial of unknown contains enough sample for at least six trials. No additional unknown will be provided! Should an unknown be spilt, a different unknown will be obtained and you will start that unknown’s analysis from the beginning.
2. Do one trial titration using between 0.1 to 0.4 g of the unknown acid. Be careful!! It is easy to dump in too much solid. Titrate as before.
3. If your initial titration volume is less than 20 mL use a little more unknown. If your titration volume was greater than 25 mL use a little less unknown. (The mass of acid and the volume of NaOH solution used in the titration are directly proportional.) Do at least two more trials.
4. For each trial, calculate the mass of your unknown acid required to neutralize one mole of hydroxide ion.
5. Using the calculated mass of acid/mole OH- values, determine the percent relative average deviation. Do additional trials if your deviation is greater than 2%.
6. Report the average grams acid/mole hydroxide neutralized for your unknown.
E. Write a formal lab report. Carefully follow all of the instructions or you will lose points.
1. The report does not need to be typed, but must be legible, neat, and secured in a folder so that it does not fall out.
2. Use only one side of each piece of paper.
3. The report must contain the following labeled sections in this order:
a) Title page: A page with only the identifying title of the experiment, your name, and the date the report is submitted.
b) Introduction: a paragraph (be concise) describing what values you have been asked to determine in the experiment, (not how to do the experiment). Include the balanced chemical equation for the reactions used in parts B and C of this experiment. No numbers should be used in this part of your report.
c) Procedure: Do not repeat the details of the procedure given in your lab book or you will lose credit. Instead, you should write several paragraphs summarizing the theory of the procedures you used in your experiment. Some of the topics these paragraphs should cover are: What is a titration? What is a buret (maybe a sketch would be useful) and how is it used? What is the “equivalence point” in an acid/base titration? How do you know when you have reached the “equivalence point?” How is the “equivalence point” different from the “end point?” What is an indicator (in general) and what is phenolphthalein (the specific indicator used in this experiment)? No data should be presented in this part of your report.
d) Data tables: Present all of your data in neat, table form. All measurements must be included.
e) Sample calculations: Show at least one complete sample calculation for each type of calculation.
f) Results: Report your unknowns’ numbers and the average values you obtained (not the relative average deviation) for each unknown. Put nothing else in this section.
g) Error discussion: Your error discussion should first define systematic and random errors. Then it should give definitions of accuracy and precision. Now make sure your error discussion answers the following questions: Which type of error, systemic or random, affects accuracy (and the grade on your unknowns)? Which type of error affects precision (and your percent relative average deviation)? What are some possible systematic errors that could occur in this experiment? What are some possible random errors that could occur in this experiment? How would each of the possible errors in this experiment affect your results? What can be done to try to minimize each type of error?